Je Tiens Aussi À Remercier Mme

Mireille Del

M Xavier

Gaona Martinez D'avoir

Mm Pascal Pernot

Gilles Montavon

Lara Mme

Duro

Lana Abou Zeid

Thibault Vital

Florian Le Dantec

Anne-Sophie Pires

Analytiques et de Réactivité des Surfaces (SEARS) au CEA à Saclay. Cette thèse a bénéficié d'un financement de l'Organisme national des déchets radioactifs et des matières fissiles enrichies en Belgique (ONDRAF-NIRAS). Si cette thèse représente un travail personnel, je profite de cette occasion pour rendre hommage et remercier toutes les personnes qui ont contribué à la réalisation de cette recherche.

Mes premiers remerciements s'adressent naturellement à mon directeur et encadrant de thèse, M. Pascal E. Reiller pour la qualité de sa direction, ses compétences hautement spécialisées et sa disponibilité qui m'ont permis de mener à bien ce projet. Je lui suis également reconnaissante pour ses écoutes continues, ses conseils précieux, ses soutiens solides, son humour habituel et ses réelles qualités humaines qui m'ont permis de travailler dans les meilleures conditions, tout au long de la période de cette thèse, y compris celle de stage d'avant.

Résumé : La formation des complexes carbonate d'uranyle(VI) est primordiale pour la compréhension de la chimie et du transport de l'uranium. Les complexes ternaires triscarbonato-uranyle d'alcalinoterreux, Mn[UO2(CO3)3] (4-2n)-(M = Mg 2+ et Ca 2+ ) mis en évidence depuis la fin des années 1990, sont susceptibles de contrôler la chimie de l'uranium -et donc sa solubilité ou ses propriétés d'adsorptiondans des environnements variés, depuis l'eau de mer jusqu'aux milieux géologiques de stockage. Les données thermodynamiques de ces complexes restent sujettes à controverses. En plus des alcalinoterreux, l'influences des cations alcalins n'est pas claires entre l'influence de la force ionique de la solution ou l'influence des ions alcalins sur la structure des complexes. Ce travail de thèse se propose de clarifier les domaines d'existence des complexes Mn[UO2(CO3)3] (4-2n)-en fonction de la force ionique en étudiant l'influence additionnelle des cations alcalins. Le choix de la technique analytique s'est porté sur la spectrofluorimétrie laser à résolution temporelle (SLRT). La quantification des mesures de luminescence des complexes Mn[UO2(CO3)3] (4-2n)-est faite pour les analyses laboratoires. La présence des complexes étudiées dans les échantillons synthétiques peut être vérifiée par les bandes caractéristiques et les temps de décroissance. Dans ce travail, les données thermodynamiques en réponse à la température et à la force ionique pour les complexes (Mg/Ca)nUO2(CO3)3 (4-2n)-en milieux NaCl et NaClO4 sont très utiles pour estimer les effets de la température et de la salinité sur la spéciation de l'U(VI), en particulier dans les dépôts de déchets nucléaires, et pour combler les lacunes dans la base de données thermodynamiques pour ces complexes étudiées.

Title : Formation of ternary alkaline earth/uranium/carbonate complexes. Thermodynamic data and their influence on uranium speciation and solubility in different geochemical environments Keywords : Uranium, Alkaline earth metal, Carbonate Abstract : The formation of carbonato complexes of uranyl(IV) is of great importance for the understanding of the chemistry and transport of uranium. The ternary alkaline-earth triscarbonatouranyl complexes, Mn[UO2(CO3)3] (4-2n)-(M = Mg 2+ and Ca 2+ ), evidence in the late 1990's, can control the uranium chemistry -and its solubility and adsorption properties -in various environments, including deep geological disposal, surface waters, and seawater. Thermodynamic data for these ternary complexes are still a matter of debate. In addition of alkaline-earths, the influence of alkaline cations is not clearly settled between the ionic strength effect and the influence on the structure of complexes. The aim of this thesis is to clarify the existence domains of Mn[UO2(CO3)3] (4-2n)-complexes as a function of ionic strength by studying the additional influence of alkali ions.

Time-resolved laser spectrofluorimetry (TRLFS), a highly sensitive technique to probe the characteristic features of uranium species, was chosen as the analytical technique in this work. The quantification of Mn[UO2(CO3)3] (4-2n)-complexes from luminescence measurements was done for laboratory analysis. The presence of the studied complexes in the synthetic samples can be verified by their characteristic bands and decay times. In this work, the thermodynamic data with respect to temperature and ionic strength for the (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes in NaCl et NaClO4 media are very useful for estimating the effects of temperature and salinity on the speciation of U(VI), especially in nuclear waste repositories, and for filling in the gaps in the thermodynamic database for these complexes.

terreux qui sont fréquemment rencontrés en teneurs importantes dans des environnements naturels. En effet, les complexes ternaires triscarbonato-uranyl d'alcalino-terreux MnUO2(CO3)3 (4-2n)-, découverts à la fin des années 1990, sont susceptibles de contrôler la chimie de l'uranium et donc sa solubilité ou ses propriétés d'adsorption en milieu naturel. Néanmoins, les données thermodynamiques de ces complexes restent sujettes à controverses. Les calculs d'équilibre thermodynamiques, permettant l'établissement de diagrammes de spéciation, à partir des constantes de complexation permet de définir les conditions analytiques de formation des complexes considérés. En plus des alcalino-terreux, l'influence des cations alcalins n'est pas claire entre l'influence de la force ionique de la solution ou l'influence des ions alcalins sur la structure des complexes.

Ce travail de thèse se propose de clarifier les constantes et fonctions thermodynamiques, ainsi que les domaines d'existence des complexes MnUO2(CO3)3 (4-2n)-en fonction de la force ionique, en étudiant l'influence additionnelle des cations alcalins. De par les informations quantitatives et structurales que l'on peut obtenir, le choix de la technique analytique s'est porté sur la spectrofluorimétrie laser à résolution temporelle (SLRT) pour les études des complexes MnUO2(CO3)3 (4-2n)-.

La SLRT est développée depuis plusieurs années au Commissariat à l'Energie Atomique et aux Energies Alternatives à Saclay, dernièrement au Laboratoire de développement Analytique Nucléaire Isotopique et Elémentaire (LANIE). Le principe de la fluorescence est fondé sur les propriétés de luminescence de l'uranium au degré d'oxydation +VI. Lorsqu'il est soumis à une excitation lumineuse, l'uranium(VI) peut émettre une fluorescence (émission de l'état singulet excité à différents niveaux vibrationnels de l'état fondamental) caractéristique de couleur allant du bleu au jaune (majoritairement verte). Pour les atomes lourds comme les lanthanides * et actinides, † la phosphorescence (émission de l'état triplet à différents niveaux vibrationnels de l'état fondamental)peut se produire en même temps que la fluorescence. Ainsi, le terme luminescence s'utilisera dans ce travail de thèse pour décrire le processus de désexcitation de l'uranium(VI). La mesure de luminescence à proprement parlé consiste à détecter le signal de luminescence émis par les complexes d'uranium étudiés et à relier l'intensité de ce signal à la variation de concentration en complexe d'uranium(VI).

Cette mesure se réalise à l'aide d'un spectrofluorimètre. La quantification des mesures de luminescence d'uranium(VI) total se fait par la méthode des ajouts dosés en milieu acide phosphorique 5% pour les analyses de laboratoire. L'utilisation d'un laser pulsé présente l'avantage d'émettre une grande énergie d'excitation pendant une durée d'impulsion très courte et ainsi d'atteindre une limite de détection basse de l'ordre de 0.1 ng L -1 (≈ 23.8 ppt) pour l'uranium. Le principe consiste à détecter le signal de luminescence de l'uranium(VI) à un instant approprié après l'excitation pulsée qui permet de gagner en sélectivité en s'affranchissant les luminescences parasites à durée de vie courte.

Dans le cas de l'étude de complexes, tels que les complexes MnUO2(CO3)3 (4-2n)-, la présence des complexes étudiées dans les échantillons synthétiques peut être vérifiée par la variations des bandes de luminescence caractéristiques et les temps de décroissance. Cependant, la SLRT est très dépendante du milieu. Certains anions sont plutôt favorables à l'émission de luminescence de l'uranium grâce à l'effet complexant, par exemple PO4 3-ou SO4 2-. Cet effet peut également donner lieu à une inhibition de la luminescence vis-à-vis de l'uranyle chez certains anions, par exemple, les anions halogènes. Outre ces derniers, les anions perchlorates ClO4 -, sous certaines conditions de concentration, possèdent la capacité complexant la moins forte à l'égard de l'uranyle, ce qui permet de diminuer la compétition entres les ligands inertes et les complexants en solution dans l'étude de complexation d'uranyle. Dans cette optique, cette thèse traite de la formation des complexes MnUO2(CO3)3 (4-2n)-dans les électrolytes de fond NaCl et NaClO4 dont le premier-l'ion chlorure-est l'ion le plus souvent trouvé en milieu naturel, iii malgré son effet d'inhibition de la luminescence, et le second qui permet d'élargir le plus que possible le domaine expérimental de force ionique, ou d'interaction avec l'ion alcalin Na + .

Le travail de thèse commence par étudier la formation de CanUO2(CO3)3 (4-2n)-en milieu NaCl dans une gamme de force ionique la plus large que possible en tenant compte de la luminescence des espèces d'uranyle tricarbonate. Notre attention s'est focalisée sur l'application de la SLRT pour acquérir les propriétés caractéristiques de luminescence des complexes d'intérêt. Ce premier aspect est très important pour le déroulé des études suivantes. Son principal objectif est de vérifier l'efficacité de la méthode par titrage simultané du pH et du Ca 2+ utilisée dans cette recherche qui est différente de la méthode classique du titrage à pH constant et concentration en Ca variable. L'introduction de coefficient de Ringböm permet de quantifier la teneur en UO2(CO3)3 4-en fonction du pH qui permet d'éviter la précipitation de calcite (CaCO3) ou schoepite (UO3:2H2O) au cours du titrage de UO2(CO3)3 4-par Ca 2+ . La limite supérieure en force ionique a été fixée à 1 mol kgw -1 à cause de l'inhibition de luminescence de l'uranium(VI) par les ions chlorure. Lorsque la force ionique est élevée, son influence sur les attractions et répulsions entre ions, ainsi que les effets qui se produisent par suite d'une interaction collisionnelle, seront importants sur l'inhibition dynamique de luminescence. Néanmoins, vu l'importance environnementale des ions chlorure, les acquisitions des spectres de luminescence des complexes CanUO2(CO3)3 (4-2n)-ont été effectuées pour la force ionique comprise entre 0,1 et 1 mol kgw -1 NaCl.

Les spectres de luminescence des espèces d'uranyle tricarbonate sont généralement résolus en cinq bandes situées à environ 466, 485, 505, 526 et 549 nm, en bon accord avec les valeurs de la littérature. Il est évident que ces positions de pics sont décalées vers le bleu (déplacement hypsochrome) comparées à celles de l'uranyle. Les profiles spectraux de luminescence et les positions caractéristiques ne montrent aucun changement au cours du titrage en forces ioniques différentes à la résolution d'un spectromètre équipé d'un réseau à 300 traits mm -1 . Pour identifier la stoechiométrie du calcium, les analyses de pente, corrigées du coefficient de Ringböm de UO2(CO3)3 4-, ont été appliquées dans ce travail. Des ajustements linéaires satisfaisants ont permis d'évaluer les constantes de formation conditionnelle de CanUO2(CO3)3 (4-2n)-à chaque valeur de force ionique d'intérêt. Dans le cadre de cette thèse, la théorie de l'interaction spécifique (SIT) a été employé pour réaliser les extrapolations des constantes de formation à dilution infinie. Les résultats de log10β°(CanUO2(CO3)3 (4-2n)-) sont cohérents avec les données précédemment acquises dans la littérature. Cependant, les coefficients d'interactions spécifiques s'avèrent très différents des données estimées par analogie-ε(UO2(CO3)2 2-,Na + ) = ε(CaUO2(CO3)3 2-,Na + )-ou d'après les hypothèses classiques de la SIT-ε(Ca2UO2(CO3)3,NaCl) = 0. Cette différence souligne que la valeur conventionnellement adoptée de ε(UO2(CO3)2 2-,Na + ) = -0.02 kgw mol -1 à ε(CaUO2(CO3)3 2-, Na + ), et la valeur nulle attribuée à ε(Ca2UO2(CO3)3, Na + ) ne sont semblent pas valable.

Comme l'inhibition de la luminescence par les ions chlorures a limité notre première étude à des valeurs de force ionique de 1 mol kgw -1 en NaCl, l'effort d'étendre cette méthode à des concentrations d'électrolyte plus élevés et d'acquérir plus de données expérimentales a été rendue nécessaire. Par conséquent, dans la deuxième partie, le milieu NaClO4 moins inhibant de la luminescence de l'uranium(VI), a été utilisé pour obtenir l'extrapolation des constantes de formation à dilution infinie et les paramètres d'interaction spécifiques des complexes CanUO2(CO3)3 (4-2n)-sur une échelle de force ionique plus large. La limite supérieure de la force ionique a été fixé à 2,46 mol kgw -1 en raison de l'effet d'inhibition observé à plus forte concentration en perchlorate. D'après les analyses de pente et les extrapolations utilisant la formule SIT, les constantes de formation log10β°(CanUO2(CO3)3 (4-2n)-) sont en excellent accord avec notre détermination précédente en milieu NaCl. Cependant, les coefficients d'interactions spécifiques sont différentes entre NaCl et NaClO4. Le changement d'anion du milieu aqueux de Cl -à ClO4 -conduit à une diminution notable des valeurs de ε pour le même complexe. Plusieurs modèles ont été testés pour décrire la tendance des données expérimentales avec une variation de ε en fonction de la force ionique qui permet de prédire les valeurs de log10β(CanUO2(CO3)3 (4- 2n)-) jusqu'à 6 mol kgw -1 en NaClO4. Les résultats suggèrent que la formation de paires d'ions entre Ca 2+ et ClO4 -ou Cl -peut exacerber la compétition pour Ca 2+ entre UO2(CO3)3 4-et ClO4 - /Cl -. Il est fort probable que les interactions ioniques soient très complexes dans ce système, notamment en concentration d'électrolyte élevée où les interactions des ions constituant le milieu ne se limitent plus à la phase binaire-les ions inertes peuvent alors intervenir dans les réactions. Des questions sont ainsi soulevées pour estimer dans quelle mesure Na + est compétitif pour se lier à UO2(CO3)3 4-et CaUO2(CO3)3 2-par rapport à Ca 2+ , et si Ca 2+ est encombré par la formation de paires d'ions avec de anions comme ClO4 -/Cl -. Les discussions, détaillées dans le Chapitre 3, portent sur la nécessité de considérer attentivement les formations des paires d'ions et l'affinité de Na + avec CanUO2(CO3)3 (4-2n)-à des forces ioniques élevées.

La troisième partie de ce document porte sur la détermination des données thermodynamiques des complexes MgnUO2(CO3)3 (4-2n)-dans les mêmes gammes de forces ioniques en NaCl et NaClO4 que pour CanUO2(CO3)3 (4-2n)-. Cette partie vise à compléter la chimie de ces complexes d'intérêt. L'utilisation de la même approche est importante pour v éliminer les incertitudes systématiques et donc pour évaluer les résultats expérimentaux avec une fiabilité maximale. Deux réseaux de 300 et 1800 traits mm -1 ont été utilisés séparément pour acquérir les spectres de luminescence de MgnUO2(CO3)3 (4-2n)-. A l'aide du réseau de 1800 traits mm -1 , nous avons pu détecter de faibles décalages des bandes caractéristiques vers le rouge (déplacement bathochrome) pour MgnUO2(CO3)3 (4-2n)-, qui n'était pas visibles avec le réseau de 300 traits mm -1 . D'après l'analyse de pente, une seule espèce MgUO2(CO3)3 2-a été identifiée. De plus, il a été observé que la variation des constantes de formation conditionnelles en fonction de la force ionique se comportent de façon similaire à CaUO2(CO3)3 2-dans les deux types d'électrolyte-NaCl et NaClO4-, ce qui se traduit par des valeurs des paramètres d'interaction ε très proches. Les résultats soulignent encore une fois l'importance des interactions ioniques à courte distance en solution concentrée. Les applications de ces données thermodynamiques des complexes MnUO2(CO3)3 (4-2n)-ont été refaites pour plusieurs cas particuliers.

Cependant, un manque de connaissances sur les paramètres thermodynamiques avec la température pour ces complexes MnUO2(CO3)3 (4-2n)-rend difficile la prédiction du comportement chimique de l'uranium dans le champ proche des dépôts de déchets radioactifs, où un gradient de température significatif s'impose au cours du temps. Donc, nous avons ensuite mesuré les constantes de formation à températures variables, et ajusté aux différents modèles d'approximation-l'équation de Van't Hoff, le modèle de capacité thermique constante, l'équation DQUANT et la modification isoélectrique-pour déduire leurs fonctions thermodynamiques-enthalpie ΔrH°m, entropie ΔrS°m, capacité thermique ΔrCp°m. En augmentant la température, l'évolution des positions des bandes spectrales avec l'augmentation de [Mg 2+ ]/[Ca 2+ ] permet de distinguer les spectres individuels des espèces concernées, notamment à des températures plus élevées, c'est-à-dire 40, 45 et 50°C, où les décalages de bande étaient évidents.

Dans ce travail de thèse, les données thermodynamiques en réponse à la température et à la force ionique pour les complexes MnUO2(CO3)3 (4-2n)-en milieux NaCl et NaClO4 permettent d'estimer les effets de la température et de la salinité sur la spéciation de l'uranium(VI), en particulier dans les sites de stockage de déchets nucléaires, et pour combler les lacunes dans les bases de données thermodynamiques pour ces complexes. En outre, les spectres caractéristiques et les temps de décroissance expérimentaux peuvent servir à l'identification de ces espèces.

Ainsi, dans la dernière partie, nous avons examiné les principales espèces dans les compositions d'eaux naturelles liées à l'activité minière de l'uranium, à la gestion des déchets radioactifs vi (RWM), et d'autres milieux liés à l'uranium naturel avec les nouvelles données thermodynamiques déterminées pour les complexes MnUO2(CO3)3 (4-2n)-. Les contours de forme ellipsoïdale pour UO2CO3(aq), CaUO2(CO3)3 2-, et Ca2UO2(CO3)3(aq) ont été déterminées à partir de la distribution des points d'échantillon dans le diagramme d'activité log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -)a(H + )) en 3D. Deux échantillons d'eau du site de Lodève ont été analysé pour leurs compositions chimiques et ensuite modélisés avec nos nouvelles données. Il est montré que la vérification des espèces dominantes pour les échantillons de Lodève est en accord avec les résultats du calcul de spéciation. On peut conclure que la détermination des ellipsoïdes pour les complexes CanUO2(CO3)3 (4-2n)-est un outil

General Introduction Context and issues

The evaluation of environmental consequences related to the extraction of uranium for nuclear power cyclefrom exploration to recyclinghas been an important consideration in any industrial project development, such as in the framework of the safety analysis of highlevel waste (HLW) repositories, uranium mining activities, remediation of uranium-polluted soils and aquifers, as well as assessment of uranium toxicity in ecosystems. From the early days of the Manhattan Project to later development of civilian nuclear power, the demand for uranium mining has skyrocketed and thus produced considerable quantities of uranium-rich sediments or mining waste in the environment. Some of these contaminated sites are sources of local uranium contamination of groundwater. Uranium also contributes to the long-term radiotoxicity of nuclear fuel after reprocessing. The importance of uranium as a pollutant does not only depend on its quantities in nuclear waste and/or contaminated sediments, but also on the assessment of associated physico-chemical reactions that determine its mobility under specific conditions. However, it is quite limited and somewhat impossible at a laboratory scale to investigate processes that take long time to be followed by experiments, such as processes at a geological time scale.

Nuclear waste repositories

Multi-barrier nuclear waste repository systems are under consideration in the countries engaged in radioactive waste management. These systems involve first a natural barrier system, which includes the host rock and its surrounding subsurface environment, then the backfill system, which consists of bentonite in some disposal concepts 1,2 or mixtures with graphite and silica to enhance thermal conductivity in other concepts, 3 finally the waste canisters, which are considered as the most rigid barrier to prevent the release of uranium up to 10000 years.

Each repository has its own conditions and barrier properties, and their characteristics are under intense study. Currently, three types of host rockclay/shale/argillite, crystalline/granite rocks, and geologic salt formationshave been most extensively considered throughout the world. Several underground research laboratories (URL) have been constructed to obtain extensive and comprehensive scientific knowledge and assess the long-term repository isolation performance of radioactive waste. Example of disposals in crystalline/granite geology are the Grimsel Test Site in Switzerland, 4 the Fennoscandian Shield in Sweden, 5 and the Beishan area in China. 6 For clay/shale/argillite as the reference host rock for methodological studies on repository of radioactive waste, several URLs have made meaningful contributions, such as Boom clay at Mol in Belgium, 7,8 Callovo-Oxfordian clay at Bure in France, 9 and Opalinus Clay at Mont-Terri in Switzerland. 10 Out of the countries that are studying clay rock as host rock, Belgium, France, and Switzerland [11][12][13] are investigating pore water chemistry with relatively lower ionic strengthse.g., from 0.02 mol kgw -1 in Boom Clay to 0.5 mol kgw -1 in Opalinus Clay -, 11 while Japan 14 and Canada 15 are studying sedimentary rocks with higher salinities in the pore waters, e.g., > 200 g L -1 in the Canadian shield. Nevertheless, most of the research on radionuclide-water interactions have been conducted at low ionic strengths. Several nuclear waste repositories are at deep salt formations where groundwaters can have ionic strengths, such as the Waste Isolation Pilot Plant (WIPP) in Carlsbad, New Mexico in the USA, 16 and the Gorleben and Morsleben sites in Northern Germany 17 where the main ions Na + , Mg 2+ , Ca 2+ , Cl - give the important ionic strength. 18 A potential radioactive waste repository has to conform to several criteria with the core requirements that the host rock as well as the backfill should be able to prevent radionuclides migration into biosphere for projected lifetime of the disposal facilities, which is typically set to 10,000 years. Properties of crystalline/granite rocks that are appreciated include low permeability (without fractures), mechanical stability, high thermal conductivity, 6 and those of clay/shale/argillite rock consist of low permeability, low diffusion coefficient and high adsorption capacity.

Importance of Mg/CanUO2(CO3)3 (4-2n)-species

The most stable chemical form of uranium in the deep geological and reducing environment is U(IV) in the form of uraninite UO2(s), which is also the major component of used-fuel pellets.

During the oxidative dissolution of UO2(s) as a result of continuing interactions with groundwater or with H2O2 formed by α-hydrolysis of H2O 19 , insoluble U(IV) can convert to soluble U(VI) resulting in the formation of uranyl hydroxides, commonly schoepite with general formula UO3:nH2O, n ≤ 2 ( n = 2 taken in this work). 20 For example, schoepite is selected as the controlling phase in early degradation of used-fuel waste packages, 21 where the role of bentonite backfill is essential because it can affect the dissolution of schoepite by the following reaction,

UO 3 :2H 2 O+2H + ⇄UO 2 2+ +3H 2 O (1)
with log10K° = 5.96 22 taken from the Giffaut et al. and Grivé et al. 23,24 synthesis methods raised some questions since the laboratory tests have shown the association of these minerals depending on the relative availabilities of Mg 2+ , Ca 2+ , Na + in solution. 35,36 The lack of information on the uranium solubility associated with these important U(VI) phases imposes limitations on further investigations into the formation of aqueous (Mg/Ca)nUO2(CO3)3 (4-2n)-species in geological context.

In addition, it was shown that the formation of CanUO2(CO3)3 (4-2n)-species can lead to higher aqueous U(VI) concentrations within bentonite because of the surface complexation reactions occurred with exchangeable and adsorption sites of uranyl. 37 For example, the desorption of U(VI) on montmorillonite, which is the main constituent of bentonite, is illustrated as follows: mon_sOUO 2 (OH) 2 -+2Ca 2+ +3HCO 3 -⇄mon_sOH+2H 2 O+ Ca 2 UO 2 (CO 3 ) 3 (4) However, out of the ongoing research in this field, the thermodynamic characters of Ca-UO2-CO3 complexes has not received enough attention, especially their complexations at varying temperatures and high ionic strengths that are common situations in heat-generating waste disposal modellings. In the scope of Nuclear Energy Agency Thermochemical Database project (NEA-TDB), * the equilibrium constants for MnUO2(CO3)3 (4-2n)-(M = Mg, Ca, Sr, Ba) proposed in literature has only been recently accepted for the reasons detailed in the critical reviews. 38,39 The authors reanalyzed the experimental values from the literature and discussed the roles played by alkaline ions in the interactions between UO2(CO3)3 4-and M 2+ (M are alkaline earth metal) of which the effects were not considered by the majority of researchers. The fact that the stability constant for the complex MUO2(CO3)3 2-(M = Mg, Ca, Sr, Ba) has a maximum for Ca 2+ might be related to the geometric constraints of Mg 2+ . 27 The extension of these results to include the structure of MnUO2(CO3)3 (4-2n)-(M = Mg, Ca, Sr, Ba) in the analysis of coordination is necessary. There are two ways in which researchers have collected their structural properties: experimental measurement by spectroscopy and molecular dynamic simulation. Bernard et al. 26 and Kelly et al. 40 confirmed the presence of CanUO2(CO3)3 (4-2n)-(n = 1, 2) by U LIII-edge EXAFS data, noting that the EXAFS signals suffers from an overlap of the calcium atoms and the distant oxygen atoms of the carbonate groups. Lee et al. 41 identified the existence of MgnUO2(CO3)3 (4-2n)-(n = 1, 2) complexes and reported their structural information by EXAFS data. Meanwhile, theoretical studies of the (magnesium/calcium) uranyl tricarbonate depend on the levels of theory, which determines the functional of the density. It demonstrates that the calculation of the aqueous ternary complexes needs to consider explicitly the counterions and * https://www.oecd-nea.org/jcms/pl_22166/thermochemical-database-tdb-project the water molecules in the hydration sphere, since the counterions are located very close to the uranyl moiety. 42,43 Besides the investigation of their fundamental structural information, studies on thermodynamic properties at high ionic strengths and at varying temperatures to substantiate the actual formation of ternary uranium carbonate species have been scarcely conducted. In particular, the influence of background electrolyte effects on the ternary species received limited attention. To the best of our knowledge, the maximum investigated electrolyte concentrations in experiments are 0.509 mol kgw -1 NaNO3 for the Mg-UO2-CO3 system 44 and 3 mol kgw -1 NaClO4 for the Ca-UO2-CO3 system, 45 but there is a large scatter in the experimental data in the latter study. Some stability constants in 0.5 M NaCl were reported for the Mg/Ca-UO2-CO3 system by converting the values determined in 0.1 M NaCl. 46 Such ionic strength (non-ideality) correction requires the choice of a model that could estimate the ion activity coefficients under the conditions other than the experimental electrolyte concentration. Out of the models proposed in literature, three models will be discussed: the Davies equation, 47 , the specific ion interaction theory (SIT, used within the framework of the NEA-TDB reviews 38,39,48,49 and also in this work) and the Pitzer model. 50 The simple Davies equation, without any undetermined parameters, applies fairly well in describing the activity coefficient variations at low ionic strength under 0.3 mol kgw -1 . The SIT and the Pitzer model are more appropriate for the extrapolation to higher ionic strength. Only the SIT model will be developed in the followingsee Grenthe, et al. 49 for detailed analysis.

The SIT, allowing the extrapolation up to 3-4 mol kgw -1 , requires the values of specific interaction parameters ε. However, the ε(Na + , MnUO2(CO3)3 (4-2n)-) (M = Mg, Ca) for the investigated system have not been selected in Grenthe et al. 39 The review noted an inappropriate application of SIT to the experimental data from the work of Dong and Brooks, 44 which resulted in an unreasonably large interaction coefficient for (Na + , MgUO2(CO3)3 2-). Nevertheless, the lack of experimental points at high ionic strength (> 0.5 mol kgw -1 ) in literature adds difficulties to obtain the SIT parameters. Furthermore, the variation of log10β in response to electrolyte concentration provides insights into the understanding of the electrolyte effects and the ligand effects, which is a particularly important aspect of ion paring formation in the sequential binding processes between Mg 2+ /Ca 2+ and UO2(CO3)3 4-. Therefore, a systematical determination of formation constants and specific ion interaction coefficients for the Mg/Ca-UO2-CO3 system is crucial in clarifying the ion interaction mechanism as well as having a primary idea of the roles of electrolyte ions by comparing the results obtained in two common background solutions -NaCl and NaClO4.

Problematic questions and goals of this study

The complexation of alkaline earth metal ions with uranyl tricarbonate moiety has been studied mainly by titration of Mg 2+ /Ca 2+ at fixed pH value in literature. 26,41,46,51 However, the occurrence of mineral precipitation can occur if the initial pH value is low (pH < 7.8) 51 resulting in the formation of schoepite at the beginning of complexation and if the pH value is too high (pH > 8.5) 27,41,46 resulting in the formation of alkaline earth carbonate phases (magnesite, calcite) at the end of complexation. Fig. 1 (a) shows the predominance diagram of Ca-UO2-CO3 complexes at 0.1 mol kgw -1 NaCl using the Thermochimie 10a database 52 with the implemented formation constants of CanUO2(CO3)3 (4-2n)-species from the work of Lee et al. 41,51 One can observe that the predominance regions of Ca-UO2-CO3 complexes evolve with pH and calcium concentration. The species distribution at [U(VI)] = 50 µmol kgw -1 for a series of solutions for which the pH values decrease and the calcium concentrations increase simultaneously were performed as preliminary calculations. Fig. 1 (b) shows that the fractions of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) could attain up to 50-85% in the middle and at the end of entire titration, respectively. Such approach-the change of pH as titration progresses-improves the degree of complexation of calcium and rules out the undesired precipitations that is better working with compared to conventional titration at fixed pH value.

Organization of this manuscript

This document is divided into six chapters.

Chapter I details the theoretical elements of thermodynamic in solution and a description of the methodology used in this work, including the calculation of Ringböm coefficient for uranyl carbonate at different pH values, the presentation of the TRLFS technique, the measurement of pH value with original/reconditioned combined glass electrodes in NaCl/NaClO4 electrolytes and the calibration of pH-meter for the range of temperature from 5 to 50°C.

Chapter II and Chapter III present the results on the formation of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) in NaCl and NaClO4, respectively, at different ionic strengths. The luminescence quenching of chloride ions have limited our study to 1 mol L -1 NaCl ionic strength. The lesser quenching perchlorate ions slightly extended the range of ionic strength to 2.2 mol L -1 NaClO4. The accurate determinations of formation constants at different ionic strengths, based on the quantitative analysis of the time-resolved luminescence spectra, allow deducing the formation constants at the ionic strength of the experiments (log10K), the extrapolations to infinite dilution giving the constant in standard conditions (log10K°), Gibbs energies of reaction (ΔrG°m) and formation (ΔfG°m), and specific ion interaction coefficients ε((Mg/Ca)nUO2(CO3)3 (4-2n)-,Na + X -). In particular, the results determined in NaCl are useful for predictive simulation of uranium simulation in a wide range of natural water compositions. 53 These chapters are mainly organized around the publications in Dalton Transactions: Shang and

Reiller 54 , and Shang, et al. 55 Chapter IV concerns the complexation constants and specific ion interaction coefficients of MgUO2(CO3)3 2-species in NaCl and NaClO4. Two gratings of different resolutions were used to acquire MgUO2(CO3)3 2-luminescence spectra. Thus far, consistent subsets of thermodynamic constants and ε values for MnUO2(CO3)3 (4-2n)-(M = Mg, Ca) complexes are now available. The formation of ternary MnUO2(CO3)3 (4-2n)-(M = Mg, Ca) complexes application in different geochemical contexts were examined. This chapter is organized around the publication in Dalton Transactions Shang and Reiller 56 , and Inorganic Chemistry Oher, et al. 43 Chapter V reports the complexation energetics of CanUO2(CO3)3 (4-2n)-and MgUO2(CO3)3 2-, based on measurements of their equilibrium constants at variable temperatures: 5-30°C for Mg 2+ and 10-50°C for Ca 2+ . The pure spectra were recovered at 50°C for CanUO2(CO3)3 (4-2n)- and 30°C for MgUO2(CO3)3 2-using multivariate curve resolution alternating least squares analysis (MCR-ALS). Different computational methods were tested to fit the temperature-dependent formation constants. This chapter is presented as a pre-print of a publication to be submitted in Dalton Transactions.

Chapter VI presents the analysis based on speciation calculation of the major species of water compositions linked to uranium mining activity, radioactive waste management (RWM), and of other more general interest related to natural uranium with the newly determined thermodynamic data for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes described in previous chapters.

Thermodynamic calculations and TRLFS analyses were also performed for two water samples collected from the Lodève site. The calculated dominant species corroborates well the experimental observations in terms of luminescence characteristics.

Chapter I. Bibliography I -1 Introduction

The aqueous chemistry of uranium refers to the various reactions taking place when uranium is contacting water. This term covers many different research directions due to its multifaceted chemical behaviour. Many chemical reactions can control the mobility of uranium in aquifers, depending on the oxidation state of the uranium and the physico-chemical conditions of the environment. The relative stability of the oxidation states of uranium depends on Eh, pH and complexing ligands. U(IV) and U(VI) are the most stable oxidation states of uranium in water.

U(IV) is generally considered as immobile in the environment as a result of its low solubility, while U(VI) can potentially migrate through aquifers and into the biosphere by e.g., complexation, polymerization and ion exchange. [1][2][3] In aqueous solution, uranium (VI) exists in the form of trans-dioxo cations (UO2 2+ ), called uranyl cations. The linear structure of the uranyl ion, i.e., the distribution of electrons between uranium and the two oxygen atoms, gives uranium an "effective" charge of 3.3 ± 0.1 distributed in the equatorial plane of the molecule (plane perpendicular to the axis formed by the O-U-O alignment). Thus, uranyl ions are considered as hard acid according to the HSAB concept (abbreviation for "hard and soft (Lewis) acids and bases", also known as the Pearson acid-base concept [4][5][6] ) and preferably form coordination with four/five/six hard base donor ligand in the equatorial plane. 7,8 The order of affinity of the uranyl ion reported for the predominant anions in natural waters is as follows: OH -> CO3 2-> HPO4 2-> SO4 2-> NO3 ->Cl -. 9 Choppin 10 has confirmed that the uranium dissolved in some groundwater is in the form of uranyl carbonate complexes.

This chapter gives some descriptions on uranyl-carbonate systems. The interpretation of experimental data of luminescence of uranyl ion complexed with carbonate and alkaline-earth metal ions require prior knowledge about:

(a) structural data and solution chemistry related to the studied ternary complexes.

(b) luminescent properties of uranium and determination of setup parameters.

(c) methodological analysis of luminescence data and extrapolation to infinite dilution.

I -2 Electronic structure of UO 2 2+

A thorough description of the distribution of electrons in the uranium-oxygen binding orbitals is necessary for understanding the coordination properties in the uranyl ion UO2 2+ . The electron configurations in the fundamental state of oxygen (Z = 8) and uranium (Z = 92) are written as 1s 2 2s 2 2p 4 and [Rn] 5f 3 6d 1 7s 2 , respectively, where [Rn] represents the electronic configuration of radon 1s 2 2s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 4f 14 5d 10 6s 2 6p 6 . The electronic configuration of U 6+ is [Rn] 5f 0 6d 0 7s 0 , and that of O 2-is 1s 2 2s 2 2p 6 . The uranyl dication is composed of two U≡O triple bonds, expressed as [O≡U≡O] 2+ . 11 The U 5f/6d orbitals and the O 2p orbitals form the strong U≡O triple bonds containing 6 valence electrons, while U 7s shell remains nonbonding in U(VI) species due to its high orbital energy. It is noteworthy that 4/2 nonbonding orbitals (f+δ,-δ, f+ϕ,-ϕ)/(d+δ,-δ) in U 5f/6d do not participate in bonding because of their asymmetric structure towards the O 2p orbitals. However, these nonbonding orbitals are available for equatorial ligands by σ and π interactions, thus can modify U≡O bonds. 12,13 For uranyl carbonate complexes, a study using the density functional theory (DFT) calculations 14 confirmed that the coordination of carbonate can bend O≡U≡O from linear to about 176° and an orbital of uranyl-σ (5f/6d) character matched to the O 2p orbitals of the carbonate oxygens receives electrons from one of the four nonbonding U 5f orbitals in excitations.

I -3 Structures of aqueous (Mg/Ca)-UO 2 -CO 3 complexes

The structural data of (Mg/Ca)-UO2-CO3 complexes have been collected using different methods: experimental measurements e.g., X-ray absorption spectroscopy, [15][16][17] X-ray powder diffraction; 18 and theoretical modelling e.g., density functional theory (DFT), 19 hybrid DFT, 20- 23 quantum mechanical/molecular mechanical (QM/MM) approach, 24,25 and classical molecular dynamics (MD) simulations. 19,26,27 According to the literature, the uranyl tricarbonate exhibits a 6-fold pattern with two of the three oxygen atoms in carbonate coordinating to the uranyl ion, though the coordination hexagon can be distorted due to an O-U-O angle < 60° to oxygen atoms of a carbonate ligand, as shown in Fig. I -1. 7,[20][21][22][23][28][29][30][31][32] However, a conversion of one carbonate from bidentate to monodentate binding that results in a 5-fold coordination in uranyl ion is assumed to be highly possible in aqueous conditions using the Quantum Mechanical Charge Field Molecular Dynamics (QMCF-MD) simulation. 24 The structural and dynamic data of uranyl carbonate are accordingly modified due to the protonation of one carbonate. The authors explained that this discrepancy comes from the explicit solvent effects in QM/MM simulation yielding more accurate results than other methods. Li et al. 27 explored the ways in which the UO2 2+ sequentially binds with three carbonates then with two calcium ions at the molecular level in pure water and in 0.1 mol L -1 NaCl using the QM/MM MD method. They concluded that the complexation of the first Ca 2+ cannot stabilize the 6-fold coordination of UO2(CO3)3 4-group, whilst the secondary one finally contributes to a symmetric structure in the perpendicular plane of uranyl moiety in pure water.

In comparison, the presence of sodium ion in NaCl medium distorts one carbonate group from bidentate to monodentate binding and destroys the symmetric structure of UO2(CO3)3 4-group.

In fact, Na + play a competitive role in binding of Ca 2+ with UO2(CO3)3 4-that renders the formation of CanUO2(CO3)3 (4-2n)-complexes more difficult in 0.1 mol L -1 NaCl than in pure water. They reported the calculated binding free energy for the first Ca 2+ ion binding with -36.9 kJ mol -1 in pure water and -20.0 kJ mol -1 in 0.1 mol L -1 NaCl, and for the second Ca 2+ ion binding with -22.3 kJ mol -1 in pure water and -14.7 kJ mol -1 in 0.1 mol L -1 NaCl.

A parallel simulation study of (Mg/Ca)nUO2(CO3)3 (4-2n)-complexation, limited to free solvent molecule environment, was conducted in our laboratory. 14 In order to reduce the computational costs and avoid undesired divergence, a charge-compensated model with aqueous counter ionsi.e. Na + , Ca Transitions from the excited triplet state T1 to the excited singlet ground state S0 are normally prohibited due to selection rules, also the reason for which the rate constants for phosphorescence emissions are several orders of magnitude lower than those for fluorescence.

Nevertheless, these rules can be relaxed due to the intense spin-orbit coupling phenomenon, especially for molecules formed from heavy atoms such as lanthanides and actinides.

Table I -1 lists the time range for different de-excitation processes. In practice, phosphorescence persists over a short period after shutting down photon supplydue to its unfavourable decay and long decay-time -, whereas fluorescence seems to stop immediately to the human eye. Indeed, the transitions most often involved in the radiative de-excitation of lanthanides and actinides are spin changes from the excited singlet state to the triplet state.

However, the duration of this process is shorter than the characteristic times of phosphorescence. Thus, it is better to use the term luminescence, including fluorescence and phosphorescence, to describe de-excitation processes, even though fluorescence constitutes the majority of emission. Generally, research on luminescence is performed with laser wavelengths ranging from ultraviolet to visible regions (250 to 700 nm). The vibrational energy levels of the ground state, when coupled with normal thermal motion, will produce a wide range of photon energies during emission. Therefore, fluorescence is practically observed as emission intensity over a band of wavelengths rather than sharp lines. Due to the difference in energy between the starting and ending levels, emission spectra can be used to identify compounds by their characteristic bands.

Stokes shift

Excitation and emission can be considered as processes where the luminescent species absorb photons of specific quantum energy and emit secondary photons in the form of fluorescence radiation. The emission spectra are generally independent of excitation wavelengths due to the very fast internal conversion processes. For most luminescent species, the spacing of vibrational energy levels is similar for excited and ground states, resulting in the emission spectrum strongly resembling the mirror image of the absorption spectrum.

The emitted photons contain generally less energy than the photons that are absorbed. Thus, the emission spectrum occurs at lower frequencies, or longer wavelengths, compared to the absorption spectrum. This phenomenon is known as the « Stokes shift ». The primary origin of Stokes shift is the energy loss in the rapid decay of excited electrons to the lower vibrational energy levels of the first excited state. In addition, luminescence emission is often accompanied by transitions to higher vibrational energy levels of the ground state, which will result in additional loss of excitation energy during vibrational relaxations to the lower energy levels of the ground state. Apart from the above-mentioned phenomena, excited state reactions, solvent orientation effects, and resonance energy transfer may also contribute to the increase in emission wavelengths.

Static and dynamic quenching

Quenching refers to the phenomenon that luminescence intensities and/or decay-times of luminescent species are reduced by interactions with other species in solution. This phenomenon can be either static or dynamic. Static quenching describes the formation of a nonluminescent species in the ground state. The efficiency of static quenching is correlated with the rate of the non-luminescent species formation and the concentration of quencher. Therefore, static quenching will result in decreased luminescence intensities because a part of luminescent species is converted to be non-luminescent, but the static quenching has no effect on the luminescence decay-times. In dynamic quenching, the formation of a non-luminescent species occurs during the de-excitation of luminescent species. This phenomenon depends on multiple factors: the viscosity of the solution, the concentration of quencher [Q], the lifetime of the excited state τ0. Thus, dynamic quenching will induce a decrease not only in luminescence intensity, but also in luminescence decay-times. This can be summarized in the Stern-Volmer relationship.

𝜙 𝜙 0 = 1 1+𝑘 𝑞 𝜏 0 [𝑄] (I-1)
where ϕ and ϕ0 are the quantum yields of luminescence in the presence and absence of the quencher of which the concentration is noted as [Q]. kq is the quencher rate coefficient.

Absorption phenomenon

Luminescent properties can be affected by different phenomena: absorption, quenching, and complexation. 33 The absorption phenomenon, containing the pre-filter effect and the post-filter effect, will be mainly addressed in this section. Usually, the pre-filter effect is the primary origin of the nonlinearity of the dependence of the fluorescence intensity on the fluorophore concentration, especially in the case of high concentration of fluorophore (high absorbance). 34 These two common effects can be both ruled by the Beer-Lambert's law.

Beer-Lambert's law

The Beer-Lambert's law establishes a relationship between light attenuation and properties of the material through which the light travels. We consider a general case where a beam of monochromic light of intensity I0 interact with a system: the optical path length is l cm, and the transmitted light intensity is noted as I. The Lambert's Law relates the intensity I to the original intensity I0 by the formula as follows:

𝐼 = 𝐼 0 • 𝑒 -𝑘𝑙 (I-2)
where k is the extinction coefficient of the material, expressed in cm -1 .

In solution, one can demonstrate that the extinction coefficient k is proportional to the concentration of solutes and depends on photon absorptivity of the attenuating species that is stated by the Beer's law,

𝑘 = 𝜀 • 𝐶 (I-3)
where ε is the wavelength-dependent molar attenuation coefficient and depends on the incident wavelength λ. C is the concentration of the attenuating species.

The Beer-Lambert's law combines the two laws and is summerized as follows:

𝐼 = 𝐼 0 • 𝑒 -𝜖𝐶𝑙 (I-4)
In spectroscopy, the absorbance A, a dimensionless parameter, relates the optical attenuation of a specific material containing attenuating species of concentration C to absorptivity ε of the species and the optical path length. The mathematical expression can be written as follows.

𝐴 = log 10 𝐼 0 𝐼 = 𝜀 • 𝐶 • 𝑙 (I-5)
With this relationship, light emission intensity can be used to quantitate the amount of the luminescing species in a sample. In practice, unknown concentrations of luminescing species are determined by quantitative measurement of luminescence intensity. However, many of the limitations of the Beer-Lambert's law can affect luminescence measurements, e.g., inner-filter effects. These effects include pre-filter and post-filter effects.

Pre-and post-filter effects

Pre-filter effect describes the phenomenon that the excitation radiation is reduced by the absorption of species and interfering impurities before the exciting beam flux enters into the detection system. Similarly, post-filter effect refers to the decrease in the emission radiation by by the absorption of luminescing species. In order to eliminate the pre-filter and post-filter effects, a common method in spectroscopic analysis is to use diluted sample solutions.

Within the framework of this research, the concentration of uranium was controlled at 100 µmol L -1 or lower, which is consistent with the conditions where no polynuclear hydroxide species form. At such low concentrations, the systems can be assumed to show an acceptable proportionality between luminescence intensity and concentrations of luminescing substances and the inner-filter effects are small enough to be neglected. 35

Temperature-dependent properties

The emission spectral properties of uranium are highly dependent on temperature because a multitude of factors that can affect luminescence emission spectra and quantum yields respond differently to temperature change. However, it can be difficult to decide which factor is dominant over others, and generally, more than one factor will simultaneously affect the luminescence. In this section, we will discuss major temperature effects on uranium luminescence in water surrounding environment. In polar solvent such as water, the molecules become more viscous at low temperatures thus the time for solvent relaxation increases accordingly. 37 The rapid decay of excited electrons from higher to the lowest vibrational energy level of the first excited state, the so-called vibrational relaxation, can also be strongly influenced by the temperature and viscosity of the solvent. In general, we assumed that these processes are completed prior to emission. However, if the solvent viscosity is large enough to prevent emission, then one expects to observe the blue-shifted emission spectrum of the unrelaxed F state (case at low temperature). If the relaxation is much faster than the emission, which occurs for fluid solvents, then one expects to observe the red-shifted emission spectrum of the relaxed R state (case at high temperature).

The emission of any

When the relaxation rate is comparable to the emission rate, one can observe an emission spectrum of intermediate state. This is illustrated schematically in Fig. I -6.

Fig. I -6 Jablonski diagram for solvent relaxation.

In addition, the quantum yield (Q) is usually suggested to be tightly related to the nonradiative decay (knr). It has been observed in numerous fluorophores that the quantum yield is low in nonpolar solvents and increases when the solvent contains a polar additive. Eq. I-6

approximately expresses the relation between the quantum yield (Q) and the non-radiative decay (knr). As the water polarity is high at low temperature, it is reasonable that the value of rate knr decreases dramatically thus results in blue-shifted and better-resolved spectra due to high quantum yield. Regarding to changes in spectral shape, some states can be observed as a function of temperature, e.g., the formation of an internal charge-transfer (ICT) states. 38 However, for the studied uranyl tricarbonate group, further research is needed to probe the complex luminescent properties of uranium.

knr = (1-Q)/T (I-6) G = Q/T (I-7)
In summary, no single theory can help interpreting the effects of environmental factors on luminescence. In reality, previously mentioned factors should all be considered to account for the uranium luminescent behaviour with temperature, while it is beyond the scope of this work to go into detail about the various mechanisms of emission.

Experimental set-up

The experimental set-up used in the present work includes a nanosecond Nd:YAG laser oscillator (neodymium-doped yttrium-aluminium garnet laser, Continuum Surelight II) which delivers high pulse energy and the peak power required for the study. This laser source operates at a repetition rate of 10 Hz with a pulse duration around 7 ns. The advantage of this type of laser source is to emit intense light beams of short duration compared to decay-times of the luminescing radionuclides, e.g., lanthanides with oxidation degree +3, notably Eu(III) [39][40][41][42] analog of Am(III) in the 5f series -, and actinides such as U(VI) and Cm(III) are the elements most often studied using laser because of their relatively high quantum yields. The wavelengths that can be emitted by a laser source (1064 nm, 532 nm, 355 nm, 266 nm) are applied depending on actual cases. In general, the fundamental output of 1064 nm is used for laser ablation or vaporization. The second and third harmonic generations (532 nm et 355 nm)

are used to pump the pulsed dye laser or an optical parametric oscillator (only 355 nm). The third and fourth harmonic generations (355 nm et 266 nm) can be used as a photolysis laser in the production of free radicals, or direct excitation of U(VI).

By tripling the wavelength frequency to 1064 nm, the laser beam emitted from our laser source is of λ = 355 nm before entering an optical parametric oscillator (OPO), which can produce an excitation wavelength range from 192 to 2750 nm. The excitation wavelength is fixed at 450 nm in the present work. The laser beam passes through the studied sample that is placed in a cuvette holder cooled or heated by a water circuit thermostatted to a specific temperature to avoid temperature-dependent fluctuations in luminescencemorithis is particularly sensible for uranium(VI). 43 The luminescence light is collected at 90° to the incident beam in order to avoid the sample absorbance as a result of isotropic phenomenon of luminescence. A joulemeter is used to monitor possible drifts in the energy during acquisition.

The output signal from the sample is then directed to a converging lens focusing the beam on a detectorthe input slit of a spectrometerwhich is composed of a monochromator equipped with 300, 600, 1800 lines mm -1 gratings in use and a charge-coupled device (CCD) camera. The camera, cooled by Peltier effect to -15°C, is controlled to convert the light signal into an electrical signal. A controller, connecting the camera, the laser source and the computer, allows the collection of luminescence spectra with varying delay time which are then processed to obtain the decay-time of the species under study.

I -5 Thermodynamics of aqueous solutions

The aim of this part is to explain the link between the conditional constants measured at different ionic strengths in experiments and the thermodynamic constants extrapolated to infinite dilution according to the chemical theory of aqueous solution.

Law of mass action

A chemical equation at equilibrium is generally written as follows:

𝑎𝐴 + 𝑏𝐵 ⇄ 𝑐𝐶 + 𝑑𝐷 (I-8)
The influence of each species included in the equation is evaluated by the chemical potential terms, conventionally denoted by the symbol μi, which corresponds to the change in energy of a system following a variation of moles numbers of that species. The Gibbs energy ∆rG is used to account for all these energy changes from different species with the following expression.

∆ 𝑟 𝐺 = ∑ 𝑛 𝑖 𝜇 𝑖 𝑖 (I-9)
The chemical potential of a specific species, dependent on temperature T and pressure P, is composed of a standard chemical potential μi° and an activity term ai (Eq. I-10). The equilibrium state corresponds to the state in which the Gibbs energy is minimal.

𝜇 𝑖 (𝑇, 𝑃) = 𝜇 𝑖 °(𝑇, 𝑃) + 𝑅𝑇𝑙𝑛 ( 𝑎 𝑖 𝑎°) (I-10)
where a° is the activity at standard state conditions, equivalent to 1 and R ≈ 8,31451

J•K -1 •mol -1 .
The equilibrium state requires that the Gibbs energy is minimal or its differential ∆rG = 0.

This makes it possible to develop the expression ∆rG:

∆ r G= ∑ n i μ i i = ∑ n i μ i °i +RTln ( a C c •a D d a A a •a B b • a °aa °b a °ca °d) =0 (I-11)
The standard chemical potentials μi° are grouped in ∆rG° with K° the thermodynamic constant in the reference state:

a C c •a D d a A a •a B b • a °aa °b a °ca °d = exp (- ∆ r G°R T ) =K° (I-12)
The infinite dilution state is conventionally chosen as the reference state, from which comes a° = 1.

Real aqueous solutions

The experiments are most often carried out at a specific ionic strength and their activities are not easily accessible. A corrective term is then used in the law of mass action to account for the deviation from ideality. The activity coefficient can be written as follows.

a i = γ i [A] i (I-13)
With the above relationship, the law of mass action in equilibrium is expressed as follows:

𝐾°= 𝑎 𝐶 𝑐 •𝑎 𝐷 𝑑 𝑎 𝐴 𝑎 •𝑎 𝐵 𝑏 = 𝛾 𝐶 𝑐 •𝛾 𝐷 𝑑 𝛾 𝐴 𝑎 •𝛾 𝐵 𝑏 • [𝐶] 𝑐 [𝐷] 𝑑 [𝐴] 𝑎 [𝐵] 𝑏 (I-14)
with the conditional equilibrium constant K defined as:

𝐾 = [𝐶] 𝑐 [𝐷] 𝑑 [𝐴] 𝑎 [𝐵] 𝑏 (I-15)
Hence, the formation constants at infinite dilution and at a specific ionic strength can be linked by the following relation:

log 10 𝐾°= log 10 ( γ C c •γ D d γ A a •γ B b ) + log 10 𝐾 = ∆ log 10 γ + log 10 𝐾 (I-16)
where the value of γi varies according to the estimation model.

Models for ionic strength correction

Activity coefficients cannot be directly experimentally accessible. Indeed, they are parameters to describe the difference between the standard state conditions (K°, also a hypothetical condition where ai = 1 et γi = 1) and the practical working state (K). Because the conditional equilibrium constant K determined at a specific ionic strength includes the coefficients γi, an estimate of K° can be achieved by measuring K at different electrolyte concentrations. The ionic strength is defined as follows:

𝐼 𝑚 = 1 2 ∑ 𝑚 𝑖 𝑧 𝑖 2 𝑖 (I-17)
with zi the charge of the species i. Several models relying on empirical and theoretical arguments have been developed to evaluate activity coefficients at different ionic strengths.

These models are in practice used to relate experimentally measured apparent constant K to equilibrium constant K° in ideal solution. 44,45 The existing models used to evaluate activity coefficients all have the Debye-Hückel equation as starting point. The simple Debye-Hückel equation was firstly proposed with a firm theoretical basis to estimate the long-range electrostatic interactions between the ions of opposed charges. 46 However, experimental data exhibited deviations from estimated values of conditional constants with this equation outside the extreme dilution range. Different empirical parameters were introduced to extend the validity range of the simple Debye-Hückel equation. 47 The Davies model, 47,48 including an ion-size parameter with an ion-association approach, was developed to calculate activity coefficients at low ionic strengths. According to the Davies formula, activity coefficients depend only on the ionic strength within a small interval of Im, but the nature of electrolyte is not taken into consideration. This model can be reduced to the simple Debye-Hückel equation at very low electrolyte concentrations since the supplementary term tends to zero with decreasing ionic strength. For this reason, the Davies equation can be applied to a slightly wider range of ionic strength than the simple Debye-Hückel equation.

The specific ion interaction formulations [49][50][51][52] are often applied in the cases of higher electrolyte solutions up to 3.5 mol kgw -1 to estimate the activity coefficients of solutes. The SIT is specially employed when few experimental data are available to fit the ionic strength correction model, as in the case of actinide complex studies. This theory has been validated in previous studies and adopted in the NEA's critical review to justify the data selectionfor uranium see Grenthe, et al. 49 , Guillaumont, et al. 51 , and Grenthe, et al. 52 . It will be detailed in the following section.

The Pitzer 45 model is applicable to a wider field than the SIT due to its more sophisticated formulation. Indeed, more adjustable parameters the model contains, more capable they are of reproducing the experimental results. In comparison with the SIT, which has only one adjustable variablethe ion interaction coefficient -, the Pitzer formula has at least three parameters, made up of two binary and one ternary terms. Nevertheless, introducing a large number of variables is a disadvantage when some of them are unknown or difficult to estimate.

In particular, the adjustment parameters are correlated with each other in the Pitzer theory, that makes the determination more complicated for solutions containing many chemical species. 45 The SIT will be used throughout this document in order to both correct the ionic strength for the experimentally measured apparent constants K and intercompare the equilibrium constants K° extrapolated from different experimental sources in literature. Using the SIT formula, we will deduced the formation constants at infinite dilution and the interaction coefficients for the studied complexes in two different background electrolytes -NaCl and NaClO4. As to the measurement of state functions of temperature, the osmotic coefficients ϕ and the water activity aH2O at different temperatures have been recalculated and used in the extrapolations with the values shown later. (cf. Table I -6)

Davies equation

The Davies 48 equation, in its general form, depends only on the ionic strength of the inert solution. The estimation of activity coefficient for solutions of ionic strength lower to 0.3 mol L -1 can be written as follows:

log 10 𝛾 𝑖 = -𝐴𝑧 𝑖 2 ( √𝐼 1+√𝐼 -0.3𝐼) (I-18)
where the quantity A has the same value for all ions and zi is the charge on a particular ion i;

The way in which the activity coefficient corrections are performed is illustrated below for the general case of a complex formation reaction. The formation of MLj complex, with j representing the number of ligand L varying from 1 to n, is expressed by Eq. I-19.

𝑀 + 𝑗𝐿 ⇄ 𝑀𝐿 𝑗 (I-19)

According to the mass action law, the conditional cumulative complexation constants are given by the following expression.

β 𝑗 = [𝑀𝐿 𝑗 ] [𝑀][𝐿] 𝑗 (I-20)
As stated in Eq. I-21, the logarithm of the formation constants log10βj° and log10βj are related by the following relationship: According to the Davies equation (Eq. I-18), the ∆log10γ can be calculated as following:

∆ log 10 𝛾 = -𝐴∆𝑧 2 ( √𝐼 1+√𝐼 -0.3𝐼) (I-22)
with ∆𝑧 2 = 𝑧(𝑀𝐿 𝑗 ) 2 -𝑧(𝑀) 2 -𝑗𝑧(𝐿) 2

Specific ion interaction theory

The estimation of activity coefficient in the SIT equation includes two terms: the first term corresponds to the Debye-Hückel equation and contains a single variable, ionic strength Im, and the second term sums the contributions of aqueous ions. Moreover, the second Im-dependent term comes from the first approximation of the virial expansion to describe the short-range, non-electrostatic interactions at higher concentrations. This term contains the specific ion interaction coefficients ε, which is considered correlated with the charge and size of the studied ions. The formulation of SIT is expressed as follows:

log 10 𝛾 = -𝑧 𝑖 2 𝐷 + ∑ (𝜀 𝑖,𝑗 𝑚 𝑗 ) 𝑖 (I-23)
with εi,j (en kgw mol -1 ) the specific ion interaction coefficient between the ions i and j counterion in solution and mj the molar concentration of ion j:

𝐷 = 𝐴(𝑇)√𝐼 1+𝑟 𝑖 𝐵(𝑇)√𝐼 (I-24)
where A(T) and B(T) are written as follows: 53

𝐴(𝑇) = 𝑒 3 8𝜋𝑙𝑛10 √ 2000𝑁 𝐴 𝑝(𝜀 0 𝜀 𝑊 𝑘 𝐵 𝑇) 3 (I-25) 𝐵(𝑇) = 𝑒√ 2000𝑁 𝐴 𝑝𝜀 0 𝜀 𝑊 𝑘 𝐵 𝑇 (I-26)
with e the electron charge, NA the Avogadro number, p the molarity-molality conversion coefficient, ε0 the vacuum permittivity, εW the relative permittivity or dielectric constant of water relative to the vacuum permittivity, kB the Boltzmann constant, T the absolute temperature (K), ri the hydrated radius of ion i, estimated as a sphere shape.

These constants are all expressed in SI units. According to the approach of Brönsted, Guggenheim and Scarchard (abbreviated B-G-S equation), 44,54,55 A(T) and B(T) are considered independent of Im, even if they depend on the constants p and εW that vary from pure water to a solution of non-zero ionic strength. These considerations support the fact that the Debye-Hückel term first term DH considers mainly long-range electrostatic interactions when evaluating deviations from ideality of activity coefficients.

At 25°C and 1 bar, the values of A and B are respectively 0.5091 L 1/2 mol -1/2 and 0.328.10 10 L 1/2 mol -1/2 m -1 . 56 In general, the product riB is an empirical choice and fixed to 1.5 at 25°C, which minimizes the variations of εi,j with ionic strength for a number of electrolytes. 57 However, the hypothesis of εi,j independence with respect to ionic strength does not seem to be always justified at low ionic strengths experimentally and theoretically. 45 Nonetheless, the term εi,jmj is negligible compared to the second term ziD in dilute concentrations. Finally, assuming that εi,j= εj,I, validated for strong electrolytes in the development of the second virial term, one can still obtain log10γi = log10γj. It is known that the extrapolated values according to the SIT equation is not always in good agreement with the best experimental data, notably at dilute electrolyte concentrations, but the differences are often too small. 49,51,52 The basic assumptions in SIT formulations have to be outlined here. First, the summations in the second term extend over all aqueous ions j. Nevertheless, the concentrations of the electrolyte ions X + and Y -are generally much larger than those of reactions species. Thus, the medium ions make the main contribution to the value of activity factors γi. For this reason, the summations are often simplified so that only ion interaction coefficients between the ionic medium ions X + and Y -and involved reactants/products are included:

log 10 𝛾 = -𝑧 𝑖 2 𝐷 + 𝜀 𝑖, 𝑋 + 𝑜𝑢 𝑌 -[𝑋𝑌] (I-27)
Secondly, the ion interaction coefficients are usually expected to be small for ions of the same charge sign due to electrostatic repulsion. Similarly, they should also be small for the electrostatic interactions between uncharged species, for which the SIT equation is reduced to only one term. One of the assumptions in SIT proposes a nil value for the interaction coefficients of uncharged complexes. However, Ciavatta tabulated a number of non-nil interaction coefficients for neutral species. 58 For organic ligand, Hummel, et al. 59 proposed ε(EdtaH4(aq),NaCl) and ε(EdtaH4(aq),KCl) values, and Fromentin and Reiller 40 proposed ε(AdipateH2(aq), NaCl) values. * The results in these studies suggested that the numerical values of interaction coefficients could be inconsistent with those estimated from the assumptions of the SIT model.

Ion interaction coefficient of a given complex can be estimated by direct determination from the variation of equilibrium constants with ionic strength or by analogy with those of the ions of the same charge and size. It is noted in the NEA-TDB reviews that one should be prepared to verify the validity of corresponding assumptions when analyzing experimental data measured at high electrolyte concentrations or with the use of high precision techniques. Especially, it may be possible that the value of ∆ε can significantly varies from one ionic medium to another.

Hence, the specific ion interaction coefficients should be specified with the medium in which the values are obtained.

Temperature influence on SIT parameters and interaction coefficients

The effect of temperature on the Debye-Hückel term requires supplementary approximations. In this study, the values of the Debye-Hückel parameter A(T) at different temperatures below 70°C are the values used in the NEA TDB reviews. 49,51,52 With regards to the values of riB, this product will be maintained at 1.5 kgw 1/2 mol -1/2 at any temperature in this study. In fact, the values of ε and riB are strongly correlated, as previously discussed in relation to the variations of interaction coefficients and parameter B(T) with temperature. [60][61][62][63][64] In other words, the variation of riB with temperature is included in the interaction coefficient, but the latter is an empirical term of which the physical signification is not well known. In the temperature interval of 0-100°C, the variation of the Debye-Hückel term DH is extremely weak, even negligible. Therefore, the values of ε and riB will be taken as constant in the determination of formation constants of MgCanUO2(CO3)3 (2n-4)-over the experimental temperaturessee Chapter V for further details.

Modified versions of the SIT equation

As stated in the preceding section, ion interaction coefficients are assumed constant within a certain range of ionic strength. However, they are not strictly constant but vary slightly with ionic strength. In general, variation of ε is often neglected for 1:1, 1:2 or 2:1 electrolytes of concentrations less than 3.5 mol kgw -1 , supported by an amount of experimental material. 55 For high-charge type electrolytes, the concentration-dependent ε should be used to calculate the * COOH-(CH2)4-COOH activity coefficient. 65 Several equations based on the SIT model have been proposed to evaluate ε as a function of ionic strength. The approach proposed by Ciavatta 58 has been used in some cases in previous studies, expressed in Eq. I-28.

ε(j, k, I m ) = ε 1 + ε 2 log 10 I m (I-28)
However, this relationship was not recommended by the NEA TDB for the reason that the interaction coefficient ε tends to negative infinity when the ionic strength tends to zero due to the introduction of logarithm in the expression of interaction coefficient. 65 Nevertheless, the complete expression of activity coefficient justifies the validity of this description at infinite dilution because the first term of (ε1 + ε2 log10Im) Im = ε1 Im + ε2 Im log10Im. tends to zero when decreasing ionic strength 52 and the second one tends also to zero knowing that

lim 𝐼 𝑚 →0
𝐼 𝑚 𝑙𝑛 𝐼 𝑚 → 0.

Bretti et al. 66,67 added a term that takes into account the formation of weak ion pairs as follows.

ε(j, k, I m ) = ε ∞ + ε 0 -ε ∞ I m +1 (I-29)
By considering association of the ion pairs with the formation constant K, Eq. I-30 can be written as the approximation:

𝜀(𝑗, 𝑘, 𝐼 𝑚 ) 𝐼 𝑚 = 𝜀𝐼 𝑚 -𝑙𝑜𝑔 10 (1 + 𝐾 𝐼 𝑚 ) (I-30)

Pitzer equation

Apart from the above-mentioned ionic strength correction approaches, some other formulations have been developed in literature to describe the measured formations constants with high precision over a large concentration range, e.g., the Pitzer and Brewer equation, 60 the Pitzer equations, [68][69][70][71][72][73] and the Baes and Mesmer equations. 74 The Pitzer equations consists of three terms involving a Debye-Hückel-like term, a term describing binary interactions and a term describing ternary interactions. Compared to the SIT equation considering only interactions between ions of opposite sign, the Pitzer equations also take into account cationcation, anion-anion and cation/anion-neutral interactions. However, the main problem in using the Pitzer equations is the determination of the concentration-dependent second virial coefficient. In addition to that, the use of the Pitzer equations requires an amount of experimental data over a sufficiently large range of ionic strength to deduce reliable values for salt-specific parameters that is often difficult to realize in experiments, especially using radioactive materials.

Temperature functionsenthalpy and entropy calculations

Evaluation of equilibrium constants from a specified temperature to any other desired temperature can be done by using the general thermodynamic relationships. Once the values of ∆rHm° or ∆rSm° at the reference temperature and their temperature dependence are available, the estimation of formation constants at other temperatures is straightforward. However, complete thermodynamic data of this kind is rarely available for complexes formations in aqueous solution because experiments only provide information about the species that are detectable in the laboratory systems. Therefore, difference approximation methods are necessary to recalculate equilibrium constants at any desired temperature.

Appearance of new species is one of the complications in the modelling of aqueous speciation at different temperatures. Thus, the modelling of such systems is supposed to be based on experimental evidence of change in speciation. Moreover, physical properties of solvent water change strongly with temperature that influence the relative amounts of charged complexes. For instance, the dielectric constant of waterthe force between the ions to separate from each otherdecreases significantly with increasing temperature -Fig. I -7 shows the variation of dielectric water as a function of temperature following the equation of Malmberg and Maryott. 75 This decrease in water dielectric constant favours the formation of low or zero charge complexes but hinders that of polynuclear and high charges complexes. The acidity of water also decreases with increasing temperature that aggregates hydrolysis of metal ions.

Fig. I -7 Variation of dielectric water as a function of temperature

In practice, the approximation models used to determine temperature dependencies of chemical equilibria are based on oversimplified electrostatic models, which take the participating ionic species as point-charges and the solvent water as a homogeneous dielectric continuum. Two models named second-and third-law methods are generally adopted in extrapolating experimental data determined at a specific temperature to different temperatures.

The third-law model is generally applied to equilibria between different phases over a large temperature range since this model requires accurately estimated heat capacities and relative enthalpies for reactants and products of pure phase. The second-law exploration model is generally used for accurate calculations over relatively small temperature ranges in aqueous environments with heat capacities estimated around the temperature of interest. In the frame of this work, we will use the second-law extrapolation to deduce the temperature functions of (Mg/Ca)nUO2(CO3)3 (2n-4)-complexes.

Second-law extrapolations

The relationship between standard molar Gibbs energy, enthalpy, and entropy at a given temperature can be expressed as follows:

∆ 𝑟 𝐻 𝑚 °= ∆ 𝑟 𝐺 𝑚 °+ 𝑇∆ 𝑟 𝑆 𝑚 °
where the enthalpy is the sum of the Gibbs free energy and the energy "dispersed" by the reaction, or more commonly as follows.

∆ 𝑟 𝐺 𝑚 °= ∆ 𝑟 𝐻 𝑚 °-𝑇∆ 𝑟 𝑆 𝑚 ° (I-31)
The temperature dependence of the Gibbs energy is expressed as follows. It is possible to relate the temperature dependence of the Gibbs energy to the entropy and the heat capacity function at the reference temperature of 298.15K, expressed in the following:

∆ 𝑟 𝐺 𝑚 °(𝑇) = ∆ 𝑟 𝐻 𝑚 °(𝑇 0 ) + ∫ ∆ 𝑟 𝐶 𝑝,𝑚 °𝑑𝑇 𝑇 𝑇 0 -𝑇 (∆ 𝑟 𝑆 𝑚 °(𝑇 0 ) + ∫ ∆ 𝑟 𝐶 𝑝,𝑚 °𝑇 𝑑𝑇 𝑇 𝑇 0 ) (I-36)
with T0 being the reference temperature.

From the relationship between the Gibbs energy and the equilibrium constants, one can write the temperature dependence of equilibrium constant as a function of the enthalpy and the heat capacity as follows:

log 10 𝐾°(𝑇) = log 10 𝐾°(𝑇 0 ) - ∆ 𝑟 𝐻 𝑚 °(𝑇 0 ) 𝑅𝑙𝑛 (10) ( 1 𝑇 - 1 𝑇 0
) - with R the gas constant (8.31451 J K -1 mol -1 ).

Equilibrium constant at any temperature T can thus be estimated with the Eq. I-37 when log10K°(T0), ∆rCp,m° and ∆rHm°(T0) (or ∆rSm°(T0)) are known. The temperature dependence of heat capacity is not available for most aqueous species. Therefore, one can make some approximations in several situations such as extrapolations over a relatively small temperature interval assuming ∆rCp,m° constant or even zero as discussed below.

Approximations based on the second-law method

For the chemical species of which the heat capacity is not available, some estimations are necessary to be made in order to deduce log10K°(T) using Eq. I-37. We will only introduce the models used in this work.

Assumption of constant enthalpy

This assumption greatly simplifies Eq. I-37 since the standard enthalpy does not vary with temperature and the heat capacity can be assumed zero at all temperatures. The reduced equation is the integrated Van't Hoff expression:

log 10 𝐾°(𝑇) = log 10 𝐾°(𝑇 0 ) - ∆ 𝑟 𝐻 𝑚 °(𝑇 0 ) 𝑅𝑙𝑛(10) ( 1 𝑇 - 1 𝑇 0 ) (I-38)
This equation is applicable to experimental systems of a single phase or multiphase at constant total pressure with no further constraint placed on the system. For the studies of temperature dependence of solubility products, the simplification in this model should be accordingly modified. Furthermore, for an extrapolation over a temperature range equal or less than 10 K, 76 the error introduced in the resulted log10K°(T) using this equation will be within its uncertainty limits.

Assumption of constant heat capacity

For a temperature range slightly larger than about 20 K, 76 we choose another approach often used in conjunction with Eq. I-37 assuming constant heat capacity with temperature. Eq. I-37 becomes log 10 𝐾°(𝑇) = log 10 𝐾°(𝑇 0 ) -

∆ 𝑟 𝐻 𝑚 °(𝑇 0 ) 𝑅𝑙𝑛 (10) ( 1 𝑇 - 1 𝑇 0
) + ∆ 𝑟 𝐶 𝑝,𝑚 °𝑅𝑙𝑛 (10) [(

𝑇 0 𝑇 ) -1 + ln ( 𝑇 𝑇 0 )] (I-39)
Accordingly, the Gibbs energy can be expressed as follows:

∆ 𝑟 𝐺 𝑚 °(𝑇) = ∆ 𝑟 𝐺 𝑚 °(𝑇 0 ) -(𝑇 -𝑇 0 )∆ 𝑟 𝑆 𝑚 °(𝑇 0 ) + ∆ 𝑟 𝐶 𝑝,𝑚 °[𝑇 -𝑇 0 -𝑇𝑙𝑛 ( 𝑇 𝑇 0 )] (I-40)
The obtained equilibrium constants are more reliable than those obtained with the constant enthalpy assumption. The assumption of constant heat capacity is useful to calculate the Gibbs energy at higher temperatures specifically in the range 273 to 473 K. It should be noted that in some cases, the heat capacity term can be neglected if the entropy and the heat capacity are comparable in magnitude because ∆rCp,m °[∆T -T ln(T/T0)] is less important than ∆rSm °∆T over the temperature range of 273 to 473 K.

Isoelectric reactions

Isoelectric reactions are defined for the chemical reactions of the following form: the sum of positive charges among the reactants equals the sum of positive charges among the products, and the same for the negative charges among the reactants and the products. The reactions of this kind can largely balance out the electrostatic contributions to the temperature dependence of enthalpy, thus the heat capacity of reaction can be assumed constant with temperature. Using the isoelectric reaction modification, the constant enthalpy estimation model is reliable over a greatly larger temperature range up to T ≤ 473 K and the constant heat capacity estimation model up to T ≤ 623 K. 76 In this study, we can try to convert the original complexation to isoelectric reactions and fit the experimental data with the constant enthalpy estimation model.

Available speciation software and thermodynamic databases

To meet the industrial requirements of radioactive waste management, several predictive transport models have been currently developed by coupling physico-chemical knowledge of uranium and computer algorithms. The use of geochemical modelling codes have advantages and disadvantages. On one hand, such calculations can reach a high precision with results reported to even 6 significant digits. Therein comes out the disadvantage of the calculations codes as the precise results can camouflage the reliability of data sources especially to a nonexpert geochemist or environmental scientist. For some species of which no experimental data exist at all, their implementing data are sometimes estimated by correlations or analogy until sufficient experimental data are available. However, the widespread use of computational codes has to make predictions in some critical societal fields of e.g., the management of nuclear waste and remediation of toxic radioactive or conventional waste plumes. Thus, a critical assessment of literature data including tracking down their original sources, and understanding the limitations of the data being used, is essential prior to calculations because missing, or badly estimated, compounds can result in completely different speciation.

The final speciation results strongly rely on the algorithms of the used software and overall of the choice of the thermodynamic database file. Nowadays, PhreeqC (pH redox equilibrium in C), 77 For safety and assessment of radioactive waste management, a variety of databases has been developed, such as the Swiss Nagra/PSI TDB, 79 the American Lawrence Livermore National Laboratory TDB, 80,81 the Japanese JNC TDB, 82 the British HATCHES, 83 and the French ThermoChimie TDB [84][85][86] originally developed by ANDRA (France, Agency for radioactive waste management). The latter database is now covering the need for ANDRA, ONDRAF-NIRAS (Belgium), and RWM (Great Britain). The above databases share the same source of main uranium data that are critically selected and reviewed by experts participating in the NEA-TDB, then published in a series of Chemical Thermochemistry books. 49,51,56,59,[87][START_REF] Gamsjäger | Chemical Thermodynamics 6. Chemical Thermodynamics of Nickel, Chemical Thermodynamics Series[END_REF][START_REF] Olin | Chemical Thermodynamics 7. Chemical Thermodynamics of Selenium[END_REF][START_REF] Brown | Chemical Thermodynamics 8. Chemical Thermodynamics of Zirconium, Chemical Thermodynamics Series[END_REF][START_REF] Rand | Chemical Thermodynamics 11. Chemical Thermodynamics of Thorium, Chemical Thermodynamics Series[END_REF] Establishing this type of compilation requires an enormous amount of critical work, which necessarily reduces the number of data compilations as well as the number of species.

The second law of thermodynamics is the key to any geochemical solubility and speciation calculations: the chemical system attempts to attain the lowest energy state, which means minimizing the Gibbs free energy. The Gibbs free energy, as well as other thermodynamic properties, depends on multiple factors, notably temperature and pressure. The standard states (25°C, 1 bar) are adopted in calculations, as in many natural systems the temperature and pressure are often fixed. For this reason, the compiled constants in database file are defined in some limiting conditions, such as infinite dilution, and then extrapolated to the specified experimental conditions. Some computer programs allow generating equilibrium constants at higher ionic strengths with ionic correction functions: the truncated Davies model, the extended The correction functions determine the maximum ionic strength to which the extrapolation is valid. For dilute to moderately concentrated solutions, activity coefficients are sufficient to perform explorations, while Pitzer's ion interaction approach is necessary in dealing with highly concentrated solutions. This model is generally preferred over the one-parameter extrapolation one and has been incorporated into Eq3/6, PhreeqC 77,78 , ToughReact, [START_REF] Zhang | presented in part at the TOUGH Symposium[END_REF] and GWB. However, the lack of Pitzer's interaction parameters for minor or trace elements and incomplete description of thermodynamic behaviour at elevated temperature for a handful of common elements constitute a barrier to accurate calculations, though efforts have been made in compiling the Pitzer parameters. [START_REF] Wolery | Pitzer Database Development: Description of the Pitzer Geochemical Thermodynamic Database data0.ypf. Appendix I in In-Drift Precipitates/Salts Model (P. Mariner)[END_REF] Apart from the ThermoChimie TDB, 84,85 which is dedicated to modelling the near field of disposal facilities for radioactive waste in the framework of Cigéo Project (Centre industriel de stockage géologique), an in-house database PRODATA TDB specifically used in the modelling of the Orano group's mining and post-mining activities [START_REF] Reiller | [END_REF] where ((UO2)m(OH)q 2m-q ), (H + ), (UO2 2+ ) are activities of hydroxy complexes, hydrogen ion and uranyl ion, respectively; m and q are the stoichiometric coefficients of hydrolysis complexes.

The formation constants as well as the existence of several uranyl hydrolysis species have been and still are discussed in the literature. To illustrate the discrepancies between published studies on uranyl hydrolysis species, Table I -3 compares two selected thermodynamic databases: one recommended by the Nuclear Energy Agency (NEA) 49,51 and the other by the International Atomic Energy Agency (IAEA) from literature reviews carried out by international expert groups. [START_REF] Fuger | The Chemical Thermodynamics of Actinide Elements and Compounds: Part 12. The Actinide Aqueous Inorganic Complexes[END_REF] The selection of Guillaumont et al. 51 is the updated version of Grenthe et al. 49 including the estimated values different from those reported in the selection of Fuger et al. [START_REF] Fuger | The Chemical Thermodynamics of Actinide Elements and Compounds: Part 12. The Actinide Aqueous Inorganic Complexes[END_REF] This is the case for (UO2)2OH 3+ and UO2(OH)2(aq) complexes, also for some other species that are not listed there. Overall, these selections agree only on the values of the formation constants of UO2OH + , (UO2)2(OH)2 2+ , (UO2)3(OH)5 + . Fuger [START_REF] Fuger | [END_REF] then discussed the values of the equilibrium constants of the two selections and noted that the data recommended by Grenthe et al. 49 and validated by the NEA were from deeper research, i.e. internal calculation one can note that Jean Fuger participated in both programs.

Table I -3 Hydrolysed uranyl species formation reactions and selected formation constants in different databases.

Reaction log10 K °

Grenthe et al. 49 Fuger et al. [START_REF] Fuger | The Chemical Thermodynamics of Actinide Elements and Compounds: Part 12. The Actinide Aqueous Inorganic Complexes[END_REF] Guillaumont et al. The data recommended by Grenthe et al. 49 for the formation constants of UO2OH + , and (UO2)3(OH)5 + have been confirmed in subsequent studies and selections. 49 and Fuger et al. [START_REF] Fuger | The Chemical Thermodynamics of Actinide Elements and Compounds: Part 12. The Actinide Aqueous Inorganic Complexes[END_REF] Choppin and Mathur 98 have proposed a value less than -12 for formation constant of UO2(OH)2 which is close to that proposed by Fuger et al. [START_REF] Fuger | The Chemical Thermodynamics of Actinide Elements and Compounds: Part 12. The Actinide Aqueous Inorganic Complexes[END_REF] Silva 101 and Torrero et al. 102 reported a value less than or equal to -11.5 for the same complex. Guillaumont et al. 51 finally selected -(12.15 ± 0.17) for UO2(OH)2 in the updated version of 2003, which was confirmed in the last update by Grenthe, et al. 52 .

The choice of values for the formation constants of the uranyl hydrolysis species remains delicate, due to the disparities between published data. The formation constants of the uranyl hydrolysis products used in this work to perform calculations of aqueous uranyl ion species distribution are taken from Guillaumont et al. 51 It can be emphasized that the recent update of Grenthe, et al. 52 does not change the core of these data.

Uranyl carbonate complexes

The aqueous carbonate system is important in analyzing environmental problems since the carbon dioxide and the alkaline earth carbonate-containing mineralse.g., calcite (CaCO3) and dolomite (CaMg(CO3)2), and to a lesser extend nesquehonite (MgCO3:3H2O), are of high abundance in nature. 103 The principal equilibrium reactions of aqueous carbonate systems are given in Eq. I-43 -Eq. I-45, 104,105 with the Henry's Law constant log10K °H describing the equilibrium between the atmospheric CO2(g) and dissolved CO2(aq) and log10K°eq representing the formation of the carbonic acid, H2CO3 due to the dissolution of carbon dioxide gas in solution.

𝐶𝑂 In an open system, the total carbonate concentration is correlated with the amount of atmospheric CO2 that dissolves into solution. 106 Thus the concentrations of HCO3 -and CO3 2- are dependent on the partial pressure of CO2(g) to which the solution is exposed. As an example, In the presence of dissolved carbonates, the formations of uranyl carbonate complexes are expressed as follows:

𝑚𝑈𝑂 2 2+ + 𝑞𝐶𝑂 3 2-= (𝑈𝑂 2 ) 𝑚 (𝐶𝑂 3 ) 𝑞 2𝑚-2𝑞 (I-48) 𝛽 𝑞.𝑚 0 = ((𝑈𝑂 2 ) 𝑚 (𝐶𝑂 3 ) 𝑞 2𝑚-2𝑞 ) (𝑈𝑂 2 2+ ) 𝑚 (𝐶𝑂 3 2-) 𝑞 (I-49)
where ((UO2)m(CO3)q 2m-2q ), (UO2 2+ ) and (CO3 2-) are activities of uranyl carbonate complexes, uranyl ion and carbonate, respectively; m and q are stoichiometric coefficients of complexation reactions.

Ringböm coefficient calculations

The concept of Ringböm coefficientalso called side-reaction coefficient or α coefficientwas first introduced by Schwarzenbach 107 to simplify equilibrium calculations in analytical chemistry. Ringböm et al. [108][109][110] later suggested that practically all occurring side reactions with the responsible interfering species should be accounted to estimate their influence on a certain main reaction.

Considering the formation of a mononuclear complex ML by a metal ion M with a ligand L, the main reaction may be affected by hydrogen or hydroxide ions in three ways: the hydrolysis side reaction of M forming M(OH)n, the side reaction of L with hydrogen ions forming HnL, and the side reactions of ML with hydrogen or hydroxide ions forming MHnL or M(OH)nL. In the usual expression of Ringböm coefficient, the sums of the concentrations of the original species and the interfering species are denoted by primed symbols. For instance, [(ML)'] denotes the sum of [M] and [MHnL] or [M(OH)nL]. Thus, the Ringböm coefficient αML of ML is defined by the expression

𝛼 𝑀𝐿 = [ (𝑀𝐿)′] [𝑀𝐿] (I-50)
It is advantageous in this study to determine the stability constants of MnUO2(CO3)3 (4-2n)-complexes in a way different from the ones commonly used in dealing with UO2(CO3)3 4-complex. The determination of these constants seems a rather complicated problem. In literature, to know the existence of MnUO2(CO3)3 (4-2n)-complexes, solutions containing a constant UO2(CO3)3 4-complex at fixed pH value were examined varying the concentration of alkaline earth metal ion M 2+ and determining the ratio of the number of metal ions M 2+ complexed to UO2(CO3)3 4-. [111][112][113][114][115] The complexation can be described with the conditional equilibrium constants Kn.1.3 in the following, nM 2+ +UO 2 (CO 3 ) 3 4-⇄ M n UO 2 (CO 3 ) 3

(4-2n)- (I-51) K n.1.3 = [M n UO 2 (CO 3 ) 3 (4-2n)-] [M 2+ ] n [UO 2 (CO 3 ) 3 4-] (I-52)
where squared brackets indicate concentration in molality (mol kgw -1 ). However, such approaches are restricted to avoid the occurrence of schoepite and magnesite/calcite that will be detailed in the following chapter. Therefore, the titration of M 2+ at varying pH values is used in this work with the need to determine the influence of side reactions at different pH values on the concentration of UO2(CO3)3 4-.

As presented in Eq. I-53, n stands for the stoichiometric number of calcium, R is the concentration ratio between complexed and non-complexed UO2(CO3)3 4-by M 2+ , α is the Ringböm constant of UO2(CO3)3 4-considering all possible side reactions at different pH values and the intercept log10Kn.1.3 refers to the conditional stepwise formation constants of MnUO2(CO3)3 (4-2n)-complexes at a given ionic strength.

log 10 R = log 10

[M n UO 2 (CO 3 ) 3 (4-2n)-] [UO 2 (CO 3 ) 3 4-]/α = log 10 K n.1.3 + nlog 10 [M 2+ ] (I-53)
From an experimental point of view, the concentration of complexed MnUO2(CO3)3 (4-2n)-is estimated based on their luminescent properties, i.e. luminescence intensity and decay time.

Thus, Eq. I-54 can be converted in the following form.

log 10 R = log 10 FI 0 (M n UO 2 (CO 3 ) 3 (4-2n)-) FI 0 (UO 2 (CO 3 ) 3 4-)/α = log 10 K n.1.3 + nlog 10 [M 2+ ] (I-54)
where FI0(MnUO2(CO3)3 (4-2n)-) is the luminescence intensity extrapolated at D = 0 of the complexed species at variable pH values, FI0(UO2(CO3)3 4-) is that of the non-complexed triscarbonatouranyl complex at pH 9.

The rationale behind the use of Ringböm coefficient species is actually comprehensible due to the fact that the proportion of UO2(CO3)3 4-in uranyl carbonate system suffers a crucial decrease when the aqueous conditions change from alkaline to neutral. atm, calculated using the PhreeqC software and the Thermochimie 10a database. 86 One can conclude that the distribution of uranyl carbonate complexes strongly depends on the pH value.

Under the present aqueous conditions, the dominant species UO2(CO3)3 4-at pH 9 is successively substituted by the less-complexed species-e.g., UO2(CO3)2 2-and UO2CO3(aq) -, the hydroxo speciese.g., UO2(OH) + and UO2(OH)2(aq)and the polynuclear speciese.g., (UO2)2CO3(OH)3 -, (UO2)3(OH)5 + …when the pH value decreases. The typical description and thermodynamic formation constant of hydrolysis-carbonation equilibria with UO2(CO3)3 4-are often written relative to UO2(CO3)3 4-and CO3 2-as follows.

m UO 2 (CO 3

) 3 4-+ n H 2 O ↔ (UO 2 ) m (CO 3 ) q (OH) n (2m-n-2q)+ +(3m-q) CO 3 2-+ n H + (I-55)
Varying pH values means that CO3 2-concentrations should be recalculated. In our case, it seems more convenient to express the complexation relative to UO2(CO3)3 4-and the partial pressure of CO2(g) because P(CO2) is maintained constant throughout the experiments. Using

Eq. I-56, the relationship between CO3 2-and CO2(g), the hydrolysis of UO2(CO3)3 4-can be expressed in CO2(g), as shown in Eq. I-57 instead.

(3m-q)CO 3 2-+ 2(3m-q)H + ↔(3m-q) CO 2 (g)+(3m-q)H 2 O (I-56) m UO 2 (CO 3 ) 3 4-+(6m-2q-n)H + ↔ (UO 2 ) m (CO 3 ) q (OH) n (2m-n-2q)+ +(3m-q)CO 2 (g)+(3m-q-n)H 2 O (I-57)
with the expression of equilibrium constant:

β" m.n.q 0 * = ((UO 2 ) m (CO 3 ) q (OH) n (2m-n-2q)+ )•P(CO 2 ) (3m-q) •a(H 2 O) (3m-q-n) (UO 2 (CO 3 ) 3 4-) m •(H + ) (6m-2q-n) (I-58)
also written as: log 10 * β°" m.n.q = log 10 * β°m .n.q -m log 10 β°1 .0.3 -(3m -q) log 10 β°0 .1.1 (I-59) with log10β°m.n.q the global formation constant of (UO2)m(CO3)q(OH)n (2m-n-2q)+ . According to the SIT approach, the relationship between the conditional formation constant log10 * β"m.n.q and log10 * β°m.n.q at infinite dilution is expressed as follows:

log 10 * β" m.n.q = log 10 * β°" m.n.q + m log 10 γ UO 2 (CO 3 ) 3 4-+ (6m -2q -n)log 10 γ H + - log 10 γ m.n.q -(3m -q -n) log 10 a(H 2 O) (I-60)
In the scope of this research, we will consider the interfering species, listed in Table I -5, in side reactions of UO2(CO3)3 4-.

Table I -5 Side reactions of UO2(CO3)3 4-and calculated stability constants of the interfering species.

Species m n q Reactions

UO2(CO3)2 2- 1 0 2 UO2(CO3)3 4-+ 2H + ⇄ UO2(CO3)2 2-+ CO2(g) + H2O log10 * β ⸰'' 1.0.2 = log10 * β ⸰ 1.0.2 -log10β ⸰ 1.0.3 -log10β ⸰ 1.0.1 log10 * β ⸰'' 1.0.2 = 12.92 ± 0.15 UO2CO3(aq) 1 0 1 UO2(CO3)3 4-+ 4H + ⇄ UO2CO3(aq) + 2CO2(g) + 2H2O log10 * β ⸰'' 1.0.1 = log10 * β ⸰ 1.0.1 -log10β ⸰ 1.0.3 -2log10β ⸰ 1.0.1 log10 * β ⸰'' 1.0.1 = 24.40 ± 0.21 (UO2)2CO3(OH)3 - 2 3 1 2UO2(CO3)3 4-+ 7H + ⇄ (UO2)2CO3(OH)3 -+ 5CO2(g) + 2H2O log10 * β ⸰'' 2.3.1 = log10 * β ⸰ 2.3.1 -log10β ⸰ 1.0.3 -5log10β ⸰ 0.1.1 log10 * β ⸰'' 2.
UO2(OH) + 1 1 0 UO2(CO3)3 4-+ 5H + ⇄ UO2(OH) + + 3CO2(g) + 2H2O log10 * β ⸰'' 1.1.0 = log10 * β ⸰ 1.1.0 -log10β ⸰ 1.0.3 -3log10β ⸰ 0.1.1 log10 * β ⸰'' 1.1.0 = 27.36 ± 0.12 UO2(OH)2(aq) 1 2 0 UO2(CO3)3 4-+ 4H + ⇄ UO2(OH)2(aq) + 3CO2(g) + H2O log10 * β ⸰'' 1.2.0 = log10 * β ⸰ 1.2.0 -log10β ⸰ 1.0.3 -3log10β ⸰ 0.1.1 log10 * β ⸰'' 1.2.0 = 20.46 ± 0.20
If CT is the analytical concentration of uranium, then a mass balance on uranium is:

C T = [UO 2 (CO 3 ) 3 4-] + [UO 2 (CO 3 ) 2 2-] + [UO 2 CO 3 ] + [UO 2 (OH) + ] + [UO 2 (OH) 2 (aq)] + 2 [(UO 2 ) 2 CO 3 (OH) 3 -] (+3 [(UO 2 ) 3 (OH) 5 + ]) (I-61)
Substituting in terms of the formation constants listed above gives:

C T =[UO 2 (CO 3 ) 3 4-]+ ∑ ∑ ∑ m β" m.n.q 0 * [UO 2 (CO 3 ) 3 4-] m-1 •[H + ] (6m-2q-n) P(CO 2 ) (3m-q) q n m (I-62) C T =[UO 2 (CO 3 ) 3 4-] (1+ β" 1.0.2 0 * [H + ] 2 P(CO 2 ) + β" 1.0.1 0 * [H + ] 4 P(CO 2 ) 2 + β" 1.1.0 [H + ] 5 P(CO 2 ) 3 0 * + β" 1.2.0 [H + ] 4 P(CO 2 ) 3 0 * ) +2 β" 2.3.1 0 * [UO 2 (CO 3 ) 3 4-] 2 •[H + ] 7 P(CO 2 ) 5 (+3 β" 3.5.0 0 * [UO 2 (CO 3 ) 3 4-] 3 •[H + ] 13 P(CO 2 ) 9 ) (I-63)
Then Eq. I-63 is a quadratic equation than can be arranged in standard form as ax 2 + bx + c = 0 with

a= 2 β" 2.3.1 0 * [H + ] 7 P(CO 2 ) 5 (I-64) b= (1+ β" 1.0.2 0 * [H + ] 2 P(CO 2 ) + β" 1.0.1 0 * [H + ] 4 P(CO 2 ) 2 ) (I-65) c= -C T (I-66) x = [UO 2 (CO 3 ) 3 4-] (I-67)
Hence, the roots of the quadratic equation, [UO2(CO3)3 4-], is expressed in Eq. I-71 with the known values of [H + ], P(CO2), CT as well as the involved formation constants * β ''

m.n.q. ∆ = b 2 -4ac (I-68) ∆ = (1+ β" 1.0.2 0 * [H + ] 2 P(CO 2 ) + β" 1.0.1 0 * [H + ] 4 P(CO 2 ) 2 ) 2 +8 β" 2.3.1 0 * [H + ] 7 P(CO 2 ) 5 C T (I-69) [UO 2 (CO 3 ) 3 4-] = -b ± √∆ 2a (I-70) [UO 2 (CO 3 ) 3 4-]= -(1+ β" 1.0.2 0 * [H + ] 2 P(CO 2 ) + β" 1.0.1 0 * [H + ] 4 P(CO 2 ) 2 )± √ (1+ β" 1.0.2 0 * [H + ] 2 P(CO 2 ) + β" 1.0.1 0 * [H + ] 4 P(CO 2 ) 2 ) 2 +8 β" 2.3.1 0 * [H + ] 7 P(CO 2 ) 5 C T (4 β" 2.3.1 0 * [H + ] 7 P(CO 2 ) 5 ) (I-71)
The values of water activity and the osmotic coefficients at investigated ionic strengths in NaCl and NaClO4 used in the calculation of Ringböm coefficients are listed in Table I -6.

I -7 Preparatory experiments

Estimation of detection limit

The detection limit of U(VI) varies from one to another in different complexing media.

Numerous studies have been carried out in various media, such as sulfuric, phosphoric, pyrophosphate, perchloric and nitric acids, 117 in order to better understand uranium luminescence and improve limit of detection. The criterion of medium selection relies upon its quenching effect on the uranyl ion, which affects both luminescence intensity and lifetime. For example, chloride is known as a classical strong quencher. 118,119 It is demonstrated that the phosphoric acid is the most convenient medium in which the detection limit of uranium remarkably decreases. Moulin et al. 120 showed that in phosphoric acid, the detection limit of uranium could be in the ng L -1 range (4.2 10 -12 mol L -1 ), even lower, but this limit could be higher for in-situ determinations of uranium in groundwater where strong chloride quencher is present. With the aim to determine uranium sensitivity of our apparatus, a series of different U(VI) concentrations solutions from 10 -8 to 10 -10 mol kgw -1 in 5% H3PO4 medium were prepared and excited at 429 nm excitation wavelength. An extension of integration time up to 1 ms was done in order to record as much signal as possible. The characteristic peaks of uranyl ions were found at 494, 516, 546, 565 nm.

Compared with the spectroscopic data obtained by Scapolan et al. 121 , the main aqueous species was identified as UO2H2PO4 + . They also reported bathochromic shifts of luminescent spectra when increasing the pH value due to the consecutive formations of UO2HPO4 and UO2PO4 - species. In . Thus, the uranium detection limit of our apparatus is estimated at 10 -10 mol kgw -1 , or 23.8 ppt in phosphoric medium. Although the uranium concentration in further experiments will not decrease to such low levels, the determination of detection limit of the apparatus is still necessary as a part of preparatory experiments.

Determination of excitation wavelength

According to the work of Meinrath et al. 122 , the highest molar absorbance was observed at 450.2 nm under the conditions where the UO2(CO3)3 4-gives the main contribution to the spectrum. Thus, the excitation wavelength is reasonable to be set at 450 nm for maximizing the efficiency.

pH Measurements

This section is dedicated to clarifying the principles of pH measurements with combined glass electrodes that will be used throughout the experiments. We will discuss the factors involved in measurements and the magnitude of effect that each factor has on measured results.

This chapter is intended to understand the thermodynamic principles involved in the pH measurements, also to relate the measured potential values to the real hydrogen concentration for electrolyte solutions of different ionic strengths and at different temperatures.

Defining pH

The pH value is a measure of acidity and defined in terms of the hydrogen ion activity as follows:

𝑝𝐻 = -𝑙𝑜𝑔 10 𝑎 𝐻 + = -𝑙𝑜𝑔 10 𝛾 𝐻 + 𝐶 𝐻 + (I-72)
where γ H + is the activity coefficient of the hydrogen ion H + and CH + is its concentration.

Since the pH glass electrode is sensitive to the actual hydrogen ion activity, the factors correlated with its definition is important. In dilute solutions, the activity coefficient can be assumed to be unity, then activity equals to concentration and the effective pH values measured by the pH meter is equal to the actual acidity. However, the difference between effective and actual concentrations increases when moving toward less dilute solutions in which the activity coefficient becomes progressively more important. In other words, samples with the same H + concentration are expected to have different measured pH values if the ionic strengths of the samples greatly vary.

In addition to ionic strength, temperature also affects 𝑎 𝐻 + , even the original speciation of the studied sample solutions. Thus, a sample pH value should not be extrapolated to another ionic strength or temperature. This means that if the pH value of a prepared solution is known at 50°C, its pH value at 25°C is not automatically known. The standard buffer solutions should be studied at the desired temperature, and the corresponding calibration curve will be naturally redefined.

Another effect that influences the pH measurements is the medium effect. This effect describes the electrostatic and chemical interactions between H + and solvent. For example, the 𝑎 𝐻 + in ethanol is much more important than in water. In practice, the exact 𝑎 𝐻 + in non-aqueous solution is usually unknown, thus the correlation between the measured value in a non-aqueous sample and the pH values obtained in the calibration with aqueous pH buffer solutions is not valid. Since all the samples in this work are prepared in aqueous solutions, the medium effect will not be discussed here.

Overall, the pH measurements of activity can be principally influenced by the ionic strength, the temperature and the type of solvent. This determines that the experimental conditions and the studied sample compositions need to be clarified when stating its measured pH value. The only valid comparison of pH values must be from the same measurement conditions and duplication of results by another person should also be performed under the originally measured conditions.

pH measuring system

A pH measuring system consists of a glass electrode, a reference and a pH meter. In the present work, a combination electrode is used to carry out all pH measurements. This electrode combines the glass and reference electrodes into a single probe. The objective of this section is to describe the functioning and parameters for glass and reference electrodes.

The combination electrode is normally constructed with the reference surrounding the glass as shown in Fig. I -13. 123,124 With the glass electrode alone, it is not sufficient to conduct the potential measurements, because the measuring circuit needs to be completed by a reference electrode. In combination electrode, the shielding is completed by the reference filling solution.

For example, saturated potassium chloride solution is often used as the filling solution to make contact with the sample solution through a junction. An Ag/AgCl internal instead of a traditional calomel internal is used in this type of electrode to make its diameter small that allows the measurements of samples in limited access containers.

Fig. I -13

The overall galvanic cell of a typical combination glass electrode.

When immersing the pH-sensitive bulb in the sample solution, an exchange equilibrium is established between the hydrogen ion in the measured solution and the ions in the glass. The reference electrode provides a stable potential against the potential from the glass electrode, between which the difference is the source of the potential displayed on the pH meter. The measured potential difference can be expressed as follows as a function of the potential differences at each interface of galvanic chain.

𝐸 𝑚𝑒𝑠 = 𝐸 𝑟𝑒𝑓 + 𝐸 𝑗 + 𝐸 𝑀 -𝐸 𝑔 (I-73)
where Eref and Eg are respectively the potential differences of Ag/AgCl in reference and glass electrodes which are stable and reproducible. EM is the potential difference between inside and outside glass membrane that is established by the hydrogen ion activity of the sample solution.

Ej is the liquid junction potential that will be discussed in detail because it varies with different reference filling solutions and inevitably causes many problems encountered in the pH measurements. According to the Nernst equation, the potential differences can be expressed as follows:

𝐸 𝑟𝑒𝑓 = 𝐸 𝐴𝑔/𝐴𝑔𝐶𝑙 °+ 𝐴 0 log 10 𝑎 𝐶𝑙 -,𝑟𝑒𝑓 (I-74) where B and C will be determined in calibration. B is needed to be measured for each specific reference filling solution if it is replaced by user. C contains the values associated with the glass electrode that cannot be modified. In the cases where the ionic strength of the sample solution is comparable to the reference solution, the liquid junction potential becomes negligible.

𝐸
However, for sample solutions of which the ionic strengths are different from that of the reference solution, liquid junction potentials are needed to be estimated individually.

Considering two solutions of the same concentration on H + but of different ionic strengths I1 and I2, the potentials measured by the same electrode are noted as Emes1 and Emes2. According to the Eq. I-78, we have:

𝐸 𝑚𝑒𝑠2 -𝐸 𝑚𝑒𝑠1 = 𝑟𝐴 0 log 10 (𝑎 𝐻 + ,𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛2 /𝑎 𝐻 + ,𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛1 ) + 𝑓(𝐼 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛2 ) -𝑓(𝐼 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛1 ) (I-81)
From the measurements of Emes2 -Emes1, the term f(Isolution2)f(Isolution1) can thus be determined with the estimation of hydrogen ion activity using the SIT formulations. (d) The stability of this electrode can be verified if the response for one pH measurement does not exceed 10 s.

Buffer solution compositions

The buffer solutions used in this work are the pH 1.679 oxalate buffer, the pH 4.005 phthalate buffer, the pH 6.865 phosphate buffer, and the pH 9.18 borax buffer (SI Analytics, Mainz, Germany). All these buffers, except the pH 1.679 oxalate buffer, are primary buffers that are recommended by the National Bureau of Standards for their greater accuracy compared with other buffers of similar pH value. The pH 1.679 buffer has a composition of 0.05 mol kgw -1 potassium hydrogen oxalate (KH3C4O8). The pH 4.005 buffer has a composition of 0.05 mol kgw -1 potassium hydrogen phthalate (KHC8H4O4). This buffer is stable to dilution and change in temperature. The pH 6.865 buffer has a composition of 0.025 mol kgw -1 potassium dihydrogen phosphate (KH2PO4) and 0.025 mol kgw -1 disodium hydrogen phosphate (Na2HPO4). This buffer can be moderately affected by dilution and temperature. The pH 9.18 buffer has a composition of 0.01 mol kgw -1 sodium tetraborate decahydrate (Na2B4O7:10H2O).

The salt in this buffer hydrolyzes to form boric acid and sodium borate. Theoretically, this buffer should be protected from atmospheric CO2 in use for maximum accuracy.

Calibration curve at different temperatures

A four-point calibration method is applied in this study with aforementioned buffer solutions. The 95% confidence limits are readily estimated by means of ordinary least squares regression, which allowed the uncertainties of pH determination below 0.05 pH units. The

Table I -9 Potential values Emes read on pH-meter for

NaClO4 solutions of Im = 0.01 to 3.5 mol kgw -1 at pcH = 2 using the reconditioned electrode of 3.5 mol kgw -1 NaClO4 reference solution and for NaCl solutions of Im = 0.12 to 3.19 mol kgw -1 at pcH = 2 using the combined-glass electrode with the filling solution of saturated KCl: liquid junction potentials ∆Emes calculated to estimate the potential differences between the solutions of a specific Im ≥ 0.

II -1 Abstract

The formation constants of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) were determined in NaCl medium at ionic strengths between 0.1 and 1 mol kgw -1 using time-resolved laser-induced luminescence spectroscopy (TRLS). Spectroluminescence titration of UO2(CO3)3 4-complex by Ca 2+ were conducted at atmospheric CO2(g) and varying pH values in order to eliminate the eventual precipitation of both schoepite (UO3:2H2O) and calcite (CaCO3) in aqueous solutions.

To identify the stoichiometry of calcium, the slope analyses corrected by the Ringböm coefficient for UO2(CO3)3 4-relative to pH and CO2(g)instead of typical expression relative to UO2 2+ and CO3 2-was applied in this work. Satisfactory linear fits assessed the conditional stepwise formation constants in the range of ionic strength employed in this work, the values of which are in good agreement with literature data at comparable ionic strengths. Extrapolations to infinite dilution were realized in the framework of the specific ion interaction theory (SIT), also providing the evaluation of the specific ion interaction coefficients. The cumulative stability constants at infinite dilution was determined to be log10°(CaUO2(CO3)3 2-) = 27.20 ± 0.04 and log10°(Ca2UO2(CO3)3(aq)) = 30.49 ± 0.05, which are in good agreement with extrapolation proposed elsewhere in literature using a different extrapolation framework. The specific ion interaction coefficients were found to be (CaUO2(CO3)3 2-,Na + ) = (0.29 ± 0.11) and (Ca2UO2(CO3)3(aq),NaCl) = (0.66 ± 0.12) kgw mol -1 . Integration of alkali metals into the ternary species may explain these positive and relatively large interaction coefficients.

Implications on the speciation of uranium in clay groundwaters, representative of radioactive waste repositories, and in seawater are discussed.

II -2 Introduction

The speciation of uranium is of great significance for a reliable prediction of its transport from disposal and storage sites of radioactive waste towards geosphere and biosphere. A variety of oxidation states and high coordination number complicates the chemical properties of uranium. 1,2 Hence, primary geochemical reactionse.g., oxido-reduction, complexation with natural and anthropogenic ligands, and adsorption onto mineralsneed to be carefully considered. 3 In particular, the ubiquitous presence of calcium and carbonate in natural water systems renders the formation and chemical behaviour of uranium in Ca-CO3 rich waters important to be elucidated. 4 The CanUO2(CO3)3 (4-2n)-complexes have been studied since their first identification in a seepage water from a uranium mine in Saxony, Germany. 5 Since then, several analytical methods were used to explore their chemical structures and behaviour, such as time-resolved laser-induced luminescence spectroscopy (TRLS), [5][6][7][8][9][10][11] anion exchange method, 12,13 potentiometryusing Ca-specific electrode -, 14 ultraviolet-visible (UV-Vis) absorption spectroscopy, 15 and extended X-ray absorption spectroscopy (EXAFS). 7,8,[16][17][18] The interaction between Ca 2+ and UO2(CO3)3 4-seems slightly more important compared to the other alkaline earth elements,. 12,13 but the typically more important concentration of calcium yield to a prevalence of CanUO2(CO3)3 (4-2n) complexes over the other ternary species under environmental conditions in general, and potentially under envisaged disposal sites of nuclear waste in particular. 2,19 Adsorption experiments have showed that the formation of CanUO2(CO3)3 (4-2n)-can influence U(VI) adsorption on minerals, 20,21 and enhance the mobility of uranium. 22,23 Therefore, the properties of CanUO2(CO3)3 (4-2n)-complexes should be studied in detail, especially their thermodynamic binding constants, which are primordial for predictive modelling.

Among the methods mentioned above, TRLS is a highly sensitive technique to probe the characteristic features of luminescent species with limited disturbance of the aqueous equilibrium. The application to uranyl systems at low concentrations has proven the suitability and appropriateness in trace analysis of uranium. 24,25 The investigations of uranium(VI)

hydrolysis have shown individual spectroscopic signatures of principle components after spectral decomposition of a series of experimental spectra as a function of pH. [26][27][28] However, uranyl complexation by carbonate has been considered faintly detectable by TRLS for a long time, which was attributable to the weak luminescence of UO2-CO3 at room temperature as a result of the quenching effects of carbonates. 29 More recently, authors have been able to collect the luminescence emission spectra of uranyl carbonate complexes and measure their decay-time, 6,30 even in a pronounced quenching environment like seawater. 17 The achievable acquisition of binary species spectrum permitted determining quantitatively the stability constants of ternary species with steadily enhanced luminescence intensity upon the association with alkaline earth cations.

Authors reported the stability constants of CanUO2(CO3)3 (4-2n)-species in different electrolyte media based on similar thermodynamic considerations. 6,8,9,14 Titrations by calcium ion, generally at a fixed pH value, starting from a specific alkaline condition, where UO2(CO3)3 4-was dominating, to where the ternary species would take precedence. However, further investigations of experimental ranges suggest that aqueous conditions in literature may result in varying degrees of oversaturation with regard to calcite (CaCO3) or schoepite (UO3:2H2O) in solution depending on the pH valuefor calcite this was clearly shown by Dong and Brooks. 12 Endrizzi and Rao performed potentiometric titration at pH 10 in 0.1 mol L -1 NaCl using Ca 2+ ion selective electrode. 14 The interfering effects of Na + and minor Mg 2+ on the stability of electrode were confirmed to be absent in their work. Nevertheless, one can notice that another principal source of bias might emerge from the potential oversaturation of calcite at the reported pH valueeven if this was clearly not observed. Similarly, in Kalmykov and Choppin, 9 onset of calcite precipitation might have occurred at pH 8.1 in 0.1 mol L -1 NaClO4 because the end of spectroluminescent titration was conducted near the saturation limit of calcite. Moreover, Lee and Yun determined the formation constants at three pH values of 7.4, 8.1, and 9 in 0.1 mol L -1 NaClO4. 6 Combined with an EDTA titration, it was found that CaUO2(CO3)3 32 It should be noted that the relevant formation constants are must-have parameters for modelling the species distribution. 19 The influence of background electrolyte effects on the ternary species received limited attention. Within the NEA-OECD selection of data, this particular aspect is dealt with using the specific interaction theory (SIT)see Guillaumont, et al. 32 for detailsthrough specific ion interaction coefficients  between a complex and the oppositely charged ion from the background electrolyte. Equivocal implications of the reported  values have been rarely assessed. To the best of our knowledge, only Dong and Brooks 13 proposed

(MgUO2(CO3)3 2-,Na + ) = -(3.0 ± 0.
3) kgw mol -1 and re-evaluated (Ca2UO2(CO3)3(aq),NaClO4) = -2.7 kgw mol -1 from Kalmykov and Choppin. 9 Other authors postulated that (MUO2(CO3)3 2-,Na + ) could be approximated in analogy to (UO2(CO3)2 2-,Na + ) = -0.02 ± 0.09 kgw mol -1 , and that (Ca2UO2(CO3)3(aq),MX) was nil as commonly stated for non-charged species. 31,32 The very high  values sometimes reported are reminiscent of theoretical calculation on the integration of alkaline ions into the structure of the CanUO2(CO3)3 (4-2n)-complexes, 33 which remains to be demonstrated experimentally.

The scarcity of systematically determined specific ion interaction coefficients for this system is a major obstacle to the accurate speciation of uranium in natural media, and to specify whether the alkaline ions are capable of integrating into the manifold structures of ternary species at high ionic strengths.

To help answering these remaining questions, the aim of this work is to determine the thermodynamic formation constants of CanUO2(CO3)3 (4-2n)-in NaCl using TRLS. The NaCl medium is chosen, instead of typical NaClO4 for inert background electrolyte, for its relevance for natural waters, particularly for radioactive waste management sites. 34,35 It is also worthy to notice that these complex are of primordial importance in the context of uranium extraction from seawater. 14 The experiments are covering a large pH range excluding any possible precipitation of solid phases by modifying the total calcium concentration. The slope analyses of spectroluminescent titrations are used to identify the respective stability fields of the species and give information about their spectroscopic properties. The stepwise formation constants measured at various ionic strengths allow proposing the thermodynamic formation constants at infinite dilution, and the specific ion interaction coefficients (CaUO2(CO3)3 2-,Na + ) and (Ca2UO2(CO3)3(aq),NaCl), in the SIT framework. Practical application of the thermodynamic constants and specific ion interaction coefficients of the complexes will then be done for the estimation of uranium speciation in water compositions representative of radioactive waste management sites, and in seawater contexts as well. The volume of each required reagent is controlled as small as possible with the aim to maintain the ionic strength at the set value. All dilution are performed by weighing.

II -3 Experimental Section

Determination of pH values and liquid junction potential

The measurements of pH values are performed at 22 °C. A combined-glass electrode (Mettler Toledo, USA) is calibrated routinely using four commercial solutions of pH 1.68, 4.01, 6.87 and 10.01 (SI Analytics, Mainz, Germany). A series of NaCl solutions of ionic strength from 0.1 to 3 mol kgw -1 acidified by HCl to pcH (= -log10[H + ]) 2 was prepared to estimate the liquid junction potential. 36 The responses of potential read on the pH-meter against ionic strengths of NaCl solution were analysed in linear regression (not shown). The multi-point calibration method is applied in this study with aforementioned buffer solutions, the electrolyte concentrations of which are known. 37 The 95 % confidence limits are readily estimated by means of ordinary least squares, 38 which allowed the uncertainty of pH determination below 0.05 pH units.

For each measuring sample, the first pH adjustment is performed after introducing NaHCO3 solution into the aqueous system. Then, the samples are let opened to the atmosphere during two hours. The values of pH would be slightly adjusted and checked for the second time, if there were any modifications. The final pH determination is carried out just before spectrum acquisition. The equilibrium with CO2(g) can be confirmed by the fact that the pH values were almost unvarying after the additions of NaHCO3(aq).

Time-Resolved Laser-Induced Luminescence Spectroscopy

A 2 mL solution of U(VI) is taken from the 15 mL polyethylene bottle and placed in a 1 cm quartz cuvette (QS101, Suprasil, Hellma Analytics). Prior to data acquisition, the cuvette is preequilibrated thermally at least 10 min in a controlled-temperature sample holder, connected to a thermostat (Peter Huber Kältemaschinenbau AG, Germany) with circulating water. A 355 nm tripled Nd:YAG laser (Surelite, Continiuum, USA) delivering about 170 mJ of energy is used as the laser excitation source, pulsed at 10 Hz with 5 ns pulse. An optical parametric oscillator system (Horizon, Continiuum, USA) is set at λex = 450 nm. The fluorescence cuvette is placed in a thermostated holder during measurements in order to minimize the temperature influence on uranyl luminescence. 39 The energy is monitored by a RJP-734 Joule-meter (Laser Probe, Inc.). The fluorescence signal is focused into a monochromator spectrometer (Acton) using a 300 lines mm -1 grating. An intensified CCD camera (Andor, UK) cooled by the Peltier effect (-

D τ i ) i (1 -exp (- W τ i )) (II -1)
where τi is fluorescence lifetime of species i, W is the gate width, FI represents fluorescence intensity at any given delay time and FI0 is that at D = 0.

Spectral data management and analysis are carried out with Origin 9.0 software (OriginLab Corporation, USA) by a specially created data processing procedure. The Levenberg-Marquardt algorithm is used in least-square fitting method. The fluorescence intensities FI obtained at various delay time are fitted to an exponential decay function in order to derive the luminescence decay-time as well as the luminescence intensity at D = 0. [40][41][42][43] It is found that the decay time of Ca-UO2-CO3 species and delay time D are in the same order of magnitude. One prominent cause is that the luminescence intensity at D = 25 ns differed greatly from FI0 and FI, thus should not be used in the analytical procedure for the sake of accuracy.

Correction for slope analysis

The methodological approach in this paper is based on the broader implications of previous studies. [5][6][7][8][9][10][11][12][13][14][15] As mentioned earlier, complexation of calcium with uranyl carbonate moiety has been investigated mainly by titration of calcium at fixed pH value. Such approaches, however, are restricted by the occurrence of schoepite and calcite. In spite of possible precipitation of minerals with actual concentrations of calcium(II) and uranium(VI), the speciation from literature improved our understanding of the Ca-UO2-CO3 system.

The predominance regions of calcium(II) and dicalcium(II) triscarbonatouranyl complexes are shown to evolve with pH and calcium concentration. More precisely, CaUO2(CO3)3 2-occurs at relatively high pH and low calcium concentration, while lower pH and higher calcium concentration favours Ca2UO2(CO3)3(aq). For ruling out the undesired precipitations, a titration of calcium at varying pH values is used in this study. The degree of complexation is estimated by the variation of fluorescence intensity as a function of the free Ca 2+ concentration. By these means, the stoichiometry of ternary species is identified by the so-called slope analysis, which provided an approach to investigate separately their properties. Preliminary calculations are performed with literature data. 12 It appears that simultaneously changed pH values and Ca 2+ concentrations allow the successive complexation to such an extent that the fractions of CaUO2(CO3)3

nCa 2+ + UO2(CO3)3 4-⇄ CanUO2(CO3)3 (4-2n)- Kn,1,3 = [ ] CanUO2(CO3)3 (4-2n)- [ ] Ca 2+ n [ ] UO2(CO3)3 4- (II -2)
where squared brackets indicate concentration on molality scale (mol kgw -1 ).

As represented in Eq. (II -4), the slope n stands for the stoichiometric number of calcium ion involved, R is the concentration ratio between complexed and non-complexed uranyl carbonate moiety by calcium considering the influence of pH, and the intercepts log10Kn.1.3 refer to the stepwise formation constants of CanUO2(CO3)3 (4-2n)-and Ca2UO2(CO3)3(aq) at a particular ionic strength value.

log10R = log10 [ ] CanUO2(CO3)3 (4-2n)- [ ] UO2(CO3)3 4-/α = log10Kn,1,3 + n log10[Ca 2+ ] (II -3)
This latter expression can be expressed as luminescence intensity,

log10R = log10 FI0( ) CanUO2(CO3)3 (4-2n)- FI0( ) UO2(CO3)3 4-/α = log10Kn.1.3 + n log10[Ca 2+ ] (II -4)
where FI0(CanUO2(CO3)3 (4-2n)-) is the fluorescence intensity extrapolated at D = 0 of the complexed species at variable pH values, FI0(UO2(CO3)3 4-) is that of the non-complexed triscarbonatouranyl moiety at pH 9, and α represents the Ringböm 44 coefficient of UO2(CO3)3 4-.

It is calculated to assess the variation of aqueous content of triscarbonatouranyl with pH (vide infra), since the denominator of R refers to the amount of UO2(CO3)3 4-at the same pH value with the numerator sample but at [Ca 2+ ] = 0 conditions.

Calculation of the Ringböm coefficient relative to UO2(CO3)3 4-

The investigated pH range varied from a circum neutral condition about pH 7.4 to a mildly alkaline condition at pH ca. 9.0. The highest pH value leads to the predominance of UO2(CO3)3 4- Many previous studies of uranyl carbonate system showed that when the solution suffered a decrease in pH, starting from a highly alkaline condition, the proportion of UO2(CO3)3 4-crucially decreased as the less-complexed speciese.g., UO2(CO3)2 2-, UO2CO3(aq) -, the hydroxo speciese.g., UO2(OH) + , UO2(OH)2(aq)and the polynuclear speciese.g., (UO2)2CO3(OH)3 -, (UO2)3(OH)5 + …increased. 28,29,[45][46][47] To reconcile these aqueous features of hydroxo and carbonato uranyl complexes, we postulated that UO2(CO3)3 The typical description and thermodynamic formation constant of hydrolysis-carbonation equilibria are often written relative to uranyl(VI) species UO2 2+ and CO3 2-as follows.

mUO2 2+ + nH2O + qCO3 2-⇄ (UO2)m(CO3)q(OH)n (2m-n-2q)+ + nH + log10°m.n.q = [ ] (UO2)m(CO3)q(OH)n (2m-n-2q)+ [ ] H + n [ ] UO2 2+ m [ ] CO3 2-q (II -5)
Varying pH values means that CO3 2-concentrations should be recalculated. In our case, it seems more convenient to express the complexation relative to UO2(CO3)3 4-and partial pressure of CO2(g) as follows instead.

mUO2(CO3)3 4-+ (6m-2q-n)H + ⇄ (UO2)m(CO3)q(OH)n (2m-n-2q)+ + (3m-q)CO2(g) + (3m-q-n)H2O log10'°m.n.q = [ ] (UO2)m(CO3)q(OH)n (2m-n-2q)+ P(CO2) (3m-q) [ ] UO2(CO3)3 4-m [ ] H + (6m-2q-n) (II -6)
The thermodynamic constant can be calculated from the selected thermodynamic data from Guillaumont, et al. 32 as follows, log10'°m.n.q = log10°m.n.qm log10°1.0.3 -(3m-q) log10°0.1.1 (II -7)

or directly form the Gibbs energy of formation, ΔfG°m.

The concentration of UO2(CO3)3 4-at each pH value can then be expressed only as a function of [H + ] and P(CO2) using the constants calculated from the known ΔfG°m taken from the NEA-OECD database, 32 as listed in Table II -1 and Table II -2, using the mass action law.

The Ringböm coefficient is introduced to quantify the fraction of U(VI) present as the UO2(CO3)3 4-complex with the mass balance on UO2(CO3)3 4-expressed in Eq. (II -8). The mathematical form is slightly simplified to a quadratic equation with one unknown as UO2(CO3)3 4-by neglecting lesser contributing terms from hydrolysed species 5 (II -8)

(UO2)n(OH)m (2n-m)+ . CT = [ ] UO2(CO3)3 4-+       1 + '1.0.2 [ ] H + P(CO2) + '1.0.1 [ ] H + 4 P(CO2) 2 + 2'2.3.1 [ ] UO2(CO3)3 4-2 [ ] H + 7 P(CO2)
where α = CT/[ ] UO2(CO3)3 4-, P(CO2) = 10 -3.5 atm, and [H + ] is converted from (H + ), measured by pH-meter. The calculation of 'm.n.l required a correction of the water activity aH 2O for each ionic strengthsee Eq. B.9 from Guillaumont, et al. 32 Ringböm coefficient was calculated for each sample at specified values of pH and ionic strength (Table II -S2 of the SI).

Specific ion interaction theory (SIT)

The SIT approach, 32 used to extrapolate the thermodynamic constants to infinite dilution, is applicable to a wider range of ionic strength compared to the Davies equation. 48 Thus, the SIT is employed to extrapolate the conditional formation constants to infinite dilution condition, and determine the specific ion interaction coefficients (CaUO2(CO3)3 2-,Na + ) and where Im is the ionic strength on the molal scale,  represents the ion interaction coefficient (kgw mol -1 ) and mk is the concentration of background electrolyte k in molality.

(Ca2UO2 (
The value of 0.507 kg 1/2 mol -1/2 is assigned to A at 22 °C. 32 The correlation between stepwise formation constants and ionic strengths can be drawn by the following relationship:

log10Kn.1.3 = log10K°n.1.3 + Δz²DH -Δ Im (II -11)
where

Δz² = z( ) CanUO2(CO3)3 (4-2n)-2 -n z(Ca 2+ ) 2 -z( ) UO2(CO3)3 4-2 (II -12)
(Δz 2 = -16 for n = 1 and Δz 2 = -24 for n = 2)

and

Δ = (CanUO2(CO3)3 (4-2n)-,Na + ) -n (Ca 2+ ,Cl -) -(UO2(CO3)3 4-,Na + ) (II -13)
The main specific ion interaction coefficients used in this work are reported in Table II -3.

II -4 Results and discussion

Evolution of luminescence spectra

The upper limit of aqueous electrolyte concentration is imposed to 1 mol kgw -1 NaCl considering the noticeable quenching effect of chloride ions. 49 Graphical representation of log10R vs. log10[Ca 2+ ] allows a straightforward examination of complexation. The aqueous calcium ion is assumed to be present as free Ca 2+ , since the analytical concentration of calcium is at least 10 times as large as that of U(VI). This simplification is attended by the fact that the amount of calcium consumed in the complexation is negligible with regard to its initial concentration. 51 The experimental data are fitted to straight lines using a least square method, 52 the slopes representing the stoichiometry of Ca 2+ ions bound to triscarbonatouranyl complex, the intercepts are assigned to the stepwise equilibrium constants at specific background electrolyte concentration. Linear relationships between log10R and log10[Ca 2+ ] are obtained with slopes changing from ca. 1 to ca. 2 as Ca 2+ is added into the solution (Fig. II -2). This also demonstrates conclusive evidence for successive formations of CaUO2(CO3)3 2-then Ca2UO2(CO3)3(aq). observed, or the rate of de-excitation process is faster than that of ligand exchange.

Nevertheless, for such a complex system containing rather complicated fluorophores, luminescence decay behaviour cannot be simply settled. Our observations are in accordance with previous studies as the same decay mode has been observed otherwise for Ca-UO2-CO3 ternary system. 6,8 However, a recent study showed clear multi-exponential decays in Mg-UO2-CO3 system, from which the characteristic decay times for principle U(VI) species were determined. 7 Numerous mutually exclusive mechanisms, especially in photochemistry, have been proposed to interpret the inconsistencies in decay mode between Ca/Mg-UO2-CO3 systems, e.g., hydrogen abstraction, 53 formation of exciplex, 54 and reversible intersystem crossing. 55 Unfortunately, these mechanisms are not likely to explain the discrepancy between experimental results, as the complexation of Ca/Mg is likely to occur outside the first coordination sphere. It seems more likely that the shorter distance of U-Mg (3.84 Å) compared to that of U-Ca (4.15 Å) could reason out the difference in fluorescent properties. 7 Several emission spectra of different compositions in binary/ternary species at Im = 0.2 and 1 mol kgw -1 are normalized to their total area in order to compare the spectroscopic changes (Fig. II -4). The percentage of species are calculated with the determined formation constants (vide post) for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq). Unambiguously, the increase in Ca 2+ concentration makes no distinctive difference in peak positions but results in clearly better resolved spectra and narrow peak widths. The similar spectral trends have been observed elsewhere. 6,7 It must be noted that the use of the 300 lines mm -1 grid favours convolution of the different components with the parameters of the optical set-up. The use of a grid with a higher resolution could help in the deconvolution of the signal and identifying spectral differences. 56

Evaluation of equilibrium constants and specific ion interaction coefficients

The graphical results in previous section provide insights into the fluorescence behaviour and complexation mechanism of the investigated system. The quenching effect of chloride is susceptible to affect the applicability of slope analysis method at high concentration, ca.

1 mol kgw -1 NaCl, given that the experimental points at log10([Ca 2+ ] mol kgw -1 ) ≥ -1.1 deviates from the linear trendnot shown in Fig. II -2. A possible explanation is probably associated with some artefacts altering the measurements of emission spectra. Because of the concentrated electrolyte, the system required a raising quantity of calcium to overcome the electrostatic force for the formation of the ternary species. This fact led to a significant increase in the actual ionic strength as a result of the contribution of [Ca 2+ ], which in turn made the impact of polarization and viscosity on fluorescence too important to be neglected. 57 However, it is yet unclear to which extent the spectral features are influenced. Notwithstanding the sparse experimental points, acceptable linear fits were retained at Im = 1 mol kgw -1 .

The calcium concentrations for predominance regions of the CanUO2(CO3)3 (4-2n)-species can be read directly from Fig. II -4. Table II -4 summarizes the detailed solution information of calcium concentrations and pH values for every species being predominantly present.

Converting back from the logarithmic form, one can infer that the predominance range of CaUO2(CO3)3 2-enlarges with ionic strength, as a result of increasing interionic attraction between Ca 2+ and Cl -.

Table II II -2. The ion interaction coefficients are evaluated to be (CaUO2(CO3)3 2-,Na + ) = (0.29 ± 0.11) and (Ca2UO2(CO3)3(aq),NaCl) = (0.66 ± 0.12) kgw mol -1 . The mean values and respective uncertainties for log10° and  are estimated using the weighed linear regression with reciprocals of the squared standard deviation (1/σi 2 ) as weights, 58 determined from the slope analysis. The uncertainty range is assessed by propagating the uncertainties from Im = 0 up to 1 mol kgw -1 . The successive calcium complexes display increasing overall formation constants, which mirrors the stability of the ternary complexes. The obtained values for log10°1.1.3 and log10°2.1.3 are consistent with literature data 12,14 within the uncertainty limits.

Discussion on the values of  and log10°

The log10° values that we have determined are comparable to other determination in the literature, specific ion interaction coefficients determined in this work are different from literature data, 9,13 and different from other -2 bearing charge complexessee Table II -S4 of the SI. It is often considered in literature that like-charged ions do share comparable  values in analogy. This was specifically the case for CaUO2(CO3) 2-, 6,14 and also for other casessee e.g. Table B-4 and B-5 in Guillaumont, et al. 32 . After an analysis from tabulated data in Na + media, the highest found negative value is (EdtaH2 2-,Na + ) = -(0.37 ± 0.14) kgw mol -1 ; 63 the highest positive value is (Cu(OH)4 2-,Na + ) = (0.19 ± 0.05) kgw mol -1 . 64 Even it has no physical meaning, a grand average of the available data yields to (X 2-,Na + ) = -(0.09 ± 0.12) kgw mol showing negative values. This clearly shows that the analogy between like-charged ions is limited.

Ciavatta 65 proposed empirical relationships to estimate missing ε values of a product from the known ε values of the other reactants. In the case of CaUO2(CO3)3 2-one can postulate a 1:1 stoichiometry from Eq. (II -2) and it comes

(CaUO2(CO3)3 2-,Na + ) = (UO2(CO3)3 4-,Na + ) + (Ca 2+ ,Cl -) 2
(CaUO2(CO3)3 2-,Na + ) = (0.07 ± 0.06) kgw mol -1 (14)
that is showing a positive value, slightly higher, yet not significantly different from the grand average. It seems though that no other type of -2 charge bearing complex can be compared with CaUO2(CO3)3 2-, and thus using other ε values in analogy seems not justified.

The case of Ca2UO2(CO3)3(aq) is more challenging as there are only a few data on globally neutral species. Even if the common hypothesis of a nil value is used very often, 31,32 Ciavatta 65 proposed  values for chloride, bromide, iodide, thiocyanate, and hydroxide complexes of Cd 2+ , Hg 2+ , and Pb 2+ in perchlorate media; Zhang and Muhammed proposed a value for Zn(OH)2(aq) complex. 66 The grand average value of (M(aq),NaClO4) = -(0.01 ± 0.06) kgw mol -1 seems to confirm the common hypothesis of a nil value.

Nevertheless, other determinations of inorganic complexes do not follow this general trend.

Ciavatta 65 determined (Hg(OH)2(aq),NaClO4) = 0.2 kgw mol -1 , and proposed that this positive value was linked to the formation of a hydrated oxocation complex HgO(H2O)x. Plyasunova, et al. 67 proposed (Co(OH)2(aq),NaClO4) = -(0.15 ± 0.10) kgw mol -1 . One can also note the Δ = (0.14 ± 0.36) kgw mol -1 Cu(OH)2(aq) in NaClO4, 68 which would yield in (Cu(OH)2(aq), NaClO4) = (0.18 ± 0.36) kgw mol -1 using  from Guillaumont, et al. 69 and Ciavatta. 65 Powell, et al. 70 proposed Δ = -(0.05 ± 0.03) kgw mol -1 for the formation of Zn(SO4)(aq) in LiClO4 media, which would yield in (Zn(SO4)(aq),LiClO4) = (0.25 ± 0.06) kgw mol -1 using  from Guillaumont, et al. 69 Otherwise, Ciavatta, 65 Hummel, et al., 63 and Fromentin and Reiller 42 proposed  values for globally neutral organic species that range from the zwitterion ethylenediaminetetracetic acid, (EdtaH4(aq),NaClO4) = -(0.29 ± 0.14) kgw mol -1 , to the adipic acid, (AdipateH2(aq),NaCl) = (0.105 ± 0.005) kgw mol -1 . It is also noteworthy that for adipic acid, the specific ion interaction coefficients proposed by Fromentin and Reiller 42 is significantly different in NaCl and NaClO4the perchlorate case being close to the nil value hypothesis. From this analysis, it seems that the hypothesis of a nil value for the  values of neutral species is not always justified.

Considering that Ca2UO2(CO3)3(aq) is a 1:2 complex regarding Eq. (II -2), and using the empirical relationship from Ciavatta 65 it comes,

(Ca2UO2(CO3)3(aq),NaCl) = (UO2(CO3)3 4-,Na + ) + 2(Ca 2+ ,Cl -) 3
(Ca2UO2(CO3)3(aq),NaCl) = (0.09 ± 0.04) kgw mol -1 (II -14)
which is positive but is not giving a value as high as the one we are obtaining.

It can be inferred that for Ca2UO2(CO3)3(aq) the common hypothesis of a nil value for ε does not seem to be justified either. It is noteworthy that Ciavatta 65 proposed another empirical relationship based on ionic radius, the values of which are not easy to retrieve for the CanUO2(CO3)3 (4-2n)-complexes.

The implications of the different thermodynamic constants and specific interaction coefficients can be seen as a function of Im through the variation of the activity coefficients (Fig. and extrapolations for CaUO2(CO3)3 2-are perfectly in agreement with the determinations in NaNO3, 12 NaCl, 14,15 and NaClO4. 15 For Ca2UO2(CO3)3(aq), there is a very good agreement with other obtained data, 8,12,15 when there is slight overestimation of Bernhard, et al., 5 and a slight underestimation of Endrizzi and Rao. 14 Up to 1 mol kgw -1 one can notice that the values that we are proposing allows a very good representation of the CanUO2(CO3)3 (4-2n)-complexes comportment. The comparison with the estimation of Kalmykov and Choppin 9 shows that our estimation is in agreement until Im = 0.5 mol kgw -1 , but do not allow understanding the comportment in more concentrated NaClO4 medium, which would require further works accounting for the discussion in Guillaumont, et al., 32 and recent theoretical modeling. 33,71 The same comment applies to the evolution of MgUO2(CO3)3 2-proposed in Dong and Brooks, 13 as the influence of ionic strength seems very high and does not compare favourably with our results on CaUO2(CO3)3 2-, which would be considered as an analogue component. Further works are clearly required to decipher these comportments, for instance in high ionic strength medium containing magnesium. 72,73

II

Structural implications for ternary species

The difference between these species, and other -2 charge bearing complexes, drew our attention to the molecular structures and actual electric charges of these species. Actually, complexation by alkali and alkaline earth metals is generally considered as weak ligand binding since the attraction between opposite-charged ions is marginally stronger than that of simple electrostatic interaction. 74 A specific term is used in SIT to describe this (partial) association, and it should be noted that the extent to which cations form complexes varies widely in solution. 75 For CanUO2(CO3)3 (4-2n)-species, the alkaline earth ions is locating in the second shell outside the inner shell for the three carbonates. Several EXAFS studies reported close bond distances for U-Ca shell, structurally exhibiting the manifold patterns of Ca-UO2-CO3 species. 7,8,[16][17][18] Lee, et al., 7 and Kelly, et al., 16 respectively reported 4.15 and 4.1 Å for U-Ca in aqueous solution, which are in agreement with a previous study that proposed the value for R(U-Ca) = 4.05 Å in the corresponding solidi.e. liebigite Ca2UO2(CO3)3(aq):10H2O. 18 These results confirmed the similar structures existing in Ca2UO2(CO3)3(aq) complex independent of its state. However, the principal limitation of EXAFS is that the interpretation of data is often interfered by the very close distances of R(U-Ca) and R(U-Odist). 7 The role of cation in complexation has been investigated through theoretical modelling. The counter-ion effects of Ca 2+ were proven to help stabilizing the fourfold negatively charged complex UO2(CO3)3 4-in aqueous environment with hybrid quantum mechanical/molecular mechanical molecular dynamics simulations. 76,77 Furthermore, classical molecular dynamics simulations have successfully described the impact of Na + on the binding between UO2(CO3)3 4-and Ca 2+ ions by simulating the effect of 0.1 mol L -1 NaCl as background electrolyte. 71 The presence of sodium provoked a distortion of carbonate denticity and thus yielded space for entering one water molecule in the primary coordination sphere in CaUO2(CO3)3 2-. As a consequence, an alteration in coordination number was induced in the equatorial plan of UO2 2+ , whereas in Ca2UO2(CO3)3(aq), sodium solely exacerbated the asymmetry of two calcium ions without changing the coordination number.

These findings provide indication for detectable influence of Na + on the structures of ternary species, which may give reasons for the positive and relatively large ε values in our study. Ionpair formation most likely accounts for the strong dependence of  on electrolyte concentration.

The results of this study further indicate that the conventionally adopted value of (UO2(CO3)2 2-,Na + ) to (CaUO2(CO3)3 2-,Na + ) by analogy is questionnable. The affinity of Na + with globally neutral complex could yield a globally positively charged-complex, which has been proposed by theoretical calculations, 33 and maybe suggested by our experimental results.

Besides, as suggested by previous investigators, the nil value of  for like-charged ions should be treated carefully, because the ion-solvent interaction is underestimated according to the assumption of SIT. 78,79 II -5 Practical applications

Practical application to the case of radioactive waste management

Once the determined thermodynamic constants of CanUO2(CO3)3 (4-2n)-complexes allow predicting aqueous speciation accurately up to Im = 1 mol kgw -1 in NaCl-based medium, the first application is the calculation of the uranium speciation in a water composition representative of a clay host rock for a geological repository for radioactive wastes. Compositions of water in equilibrium with a clay rock, which can be the host rock of a radioactive waste repository, were proposed in the Belgian context of the Boom clay, 34 and in the French context of the Callovo-Oxfordian clay. 35 Using thermodynamic data from Guillaumont, et al., 32 if one considers that uranium is mostly under its +IV redox state under these conditions, the solid phase controlling the solubility is either UO2(am,hyd) for the most soluble one, or UO2(cr) for the most insoluble one. The very low solubility of these two phases, and the concentration of carbonate in these systems, prevents the formation of polynuclear species.

Up to now, the available data in the framework of the SIT 32 are not sufficiently complete to perform a calculation in a real situation. It would require the addition of thermodynamic data from other sources. Hence, we have used the Thermochimie 9b databasethe sit.dat database fileprovided with PhreeqC, 61,62 which contains the data from Guillaumont, et al. 32 for uranium, to perform the theoretical speciation calculationthe selection of data can be found on the web pages. 80 In this version of the database, CanUO2(CO3)3 (4-2n)-complexes from Dong and Brooks 12 are present, with no ε values, which mean that the ε values are nil. The Mg 2+ , Sr 2+ , and Ba 2+ complexes are not implemented. 7,10,12,13 We have made the choice to implement the Mg 2+ and Sr 2+ complexes from Dong and Brooks 12 as an estimation of their potential importance in the input PhreeqC file with no ε values.

The redox potentials -EhSHE = -274 mV, t = 16°C, pH 8.5 for the Mol reference water in de Craen, et al. 34 , and EhSHE = -163 mV at pH 7.1 for solution A in Gaucher, et al. 35 are fixed by a partial pressure of dioxygen calculated from the EhSHE and temperature of the water -P(O2) = 10 -68.67 and 10 -65.83 atm for de Craen, et al. 34 and Gaucher, et al. 35 , respectively.

The solubility and aqueous speciation of uranium between pH 7 and 9 for the two water

Practical application to the case of uranium speciation in seawater

Another application is the distribution of ternary species as well as the effects of solvent ions in seawater. Different compositions of seawater were proposed in the literature. Large amounts of calcium and magnesium led to the predominance of MnUO2(CO3)3 (4-2n)-complexes after theoretical speciation calculations. 7,14,17 Knowing the particular dependence observed for the MgUO2(CO3)3 2-complex as a function of NaClO4 ionic strength, 13 and the lack of information on the corresponding comportment in NaCl, we will only focus our discussion on CanUO2(CO3)3 (4-2n)-complexes in the following.

Authors proposed theoretical speciation calculations using either the data from Dong and

Brooks 12 in an average Mediterranean seawater, 17 or their own determinations using simplified seawater compositions. 7,14 The predominance of either the neutral Ca2UO2(CO3)3(aq) complex, 14,17 or the -2 charge bearing CaUO2(CO3)3 2-complex, 7 was obtained. However, the water composition is relatively unbalancedelectrical balance is 4.15 10 -2 eq using PhreeqC and the activity correction is not clearly specified, but guessed as done through a simplified Debye-Hückel approximationwhich is not applicable in seawaterin the former case; 17 in the latter case, 7,14 the water compositionsimplified from Millero, et al. 81 is equilibrated and SIT with the nil hypothesis for the ε values, which seems not justified (vide ante), was used.

It seems awkward to assess which complex dominates the uranium speciation in seawater.

Millero, et al. proposed the composition of a standard seawater, 81 which can be used to perform a similar exercise, applying the thermodynamic constants (Table II -1) and specific ion interaction coefficients (Table II -3) determined in this workthe UO2CO3F -complex, 32 absent from Thermochimie 9b database, is implemented. We are proposing the theoretical speciation calculation in Fig. II -9, where CaUO2(CO3)3 2-is dominating the speciation at pH > 7, and Ca2UO2(CO3)3(aq) is a significant secondary speciesin this example, the partial pressure of CO2(g) is calculated from the total carbonate concentration in Millero, et al., 81 i.e.

10 -3.24 atm. This speciation is reminiscent of the one proposed by Lee, et al. 7 , but with a slightly less important proportion of the Ca2UO2(CO3)3(aq) complex. It is noteworthy that dolomite 24 atm from the total carbonate concentration, and using the thermodynamic data from the Thermochimie 9b database, 80 correcting the CanUO2(CO3)3 (4-2n)-data with thermodynamic constants values (Table II -1) and implementing the specific ion interaction coefficients (Table II -3) proposed in this work.

II -6 Conclusion

The thermodynamic constants log10° and Gibbs energies of formation ΔfG°m of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) have been extrapolated to infinite dilution from TRLS experiments in NaCl medium from Im = 0.1 up to 1 mol kgw -1 . The characteristics of the luminescence spectra are strikingly similar for all species, with gradually narrowed peak widths with complexation of calcium. The pH-dependent slope analysis was necessarily corrected with the Ringböm coefficient relative to UO2(CO3)3 4-complex according to the experimental methodology. As expected, no precipitation of schoepite or calcite was observed during the experiments. The stability constants are consistent with previously reported data. Conversely, the values of the specific ion interaction coefficients (CaUO2(CO3)3 2-,Na + ) and (Ca2UO2(CO3)3(aq),NaCl) are strongly positive and markedly different from the previous estimation by analogy from like-charged ions, which may suggest a high affinity of Na + with CanUO2(CO3)3 (4-2n)-species in solution. Similar studies at higher ionic strength and other background electrolyte such as sodium perchlorate media, and comparable studies on Mg 2+ are highly desirable.

II -8 Supporting Information

This supporting information contains 7 Figures and 4 Tables II -1, and those for CanUO2(CO3)3 4-2n taken from the work of Lee and Yun. 1 Other aqueous information is in Table II -S1.

III -1 Abstract

The stability constants of ternary calcium uranyl tricarbonate complexes, CaUO2(CO3)3

III -2 Introduction

The first identification of CanUO2(CO3)3 (4-2n)-species was published in 1996 by Bernhard, et al. 1 in a report of seepage water analyses from a uranium mine tailing in Saxony, Germany.

Since then, there has been a surge of interests in the chemical properties of ternary complexes of the type MnUO2(CO3)3 (4-2n)-(M = alkaline earth metal elements, i.e. Mg 2+ , Ca 2+ , Sr 2+ , Ba 2+ , n = {1;2}), their chemical structures and behaviour in solution. Various analytical methods and techniques have been employed, such as time-resolved laser-induced luminescence spectroscopy (TRLS) [1][2][3][4][5][6][7][8] , anion exchange method 9,10 , calcium(II)-selective potentiometry 11 , ultraviolet-visible (UV-vis) absorption spectroscopy 12 and extended X-ray absorption spectroscopy (EXAFS) 3,7,[13][14][15] . Results from earlier studies demonstrate the necessity and the importance of clarifying physical and chemical characteristics of these ternary complexes, as they have been proven influential on the U(IV)/U(VI) bio-oxidation/reduction 16,17 and U(VI) adsorption-absorption on minerals. [18][19][20][21][22][23] Moreover, the CanUO2(CO3)3 (4-2n)-complexes have been found as principal species in Ca-CO3 rich waters-e.g., in the vadose zone at Hanford site, USA [24][25][26][27] and from drilled wells in Southern Finland 22 -and proposed in other instances. 28 Among the research perspectives mentioned above, experimental determination of thermodynamic constants of Ca-UO2-CO3 complexes draws our major attention. For predictive modelling of species distribution, particularly in the field of radioactive waste management, the formation constants are must-have parameters, as well as specific ion interaction coefficients in the framework of the specific ion interaction theory (SIT). 29 To date, none of the reported formation constants have been selected by the Thermochemical Database project of the NEA-OECD-the Nuclear Energy Agency from the Organization for the Economic Co-operation and Development (https://www.oecd-nea.org/jcms/pl_20079/tdb-projectpublications-chemical-thermodynamics-series). 29,30 One of the main reasons is the difficulty to experimentally and accurately investigate the stability domains of such ternary species, which has led to differences in the complex formation constants. In past studies, the conventional titration of triscarbonatouranyl complex (UO2(CO3)3 4-) by Ca 2+ at fixed pH values has been widely used, 2,9,11 starting from alkaline conditionswhere UO2(CO3)3 4-predominates. The resulting analyses were based on the measures of complexation degree with increasing calcium concentration. However, the principal source of bias might emerge from the different extents of oversaturation of schoepite (UO3•nH2O) and calcite (CaCO3) depending on the pH value.

In our previous study, 8 we used a spectroluminescence titration at varying pH values in order to circumvent the possible schoepite and calcite precipitation. Satisfactory linear fits in the slope analysis proved the feasibility of the experimental methodology. The successful collection of luminescence spectra and decay-time measurements of the calciumtriscarbonato-uranyl species in the investigated range of Im = 0.1 to 1 mol kgw -1 NaCl (hereafter noted m), have demonstrated the high performance of TRLS technique and implied further applications. Nevertheless, the quenching effect of chloride ion has prevented the use of ionic strength media above 1 m. The obtained specific ion interaction coefficients ε(Na + ,CaUO2(CO3)3 2-) and ε(Ca2UO2(CO3)3(aq),NaCl) were both surprisingly positivei.e. 0.29 and 0.66 kgw mol -1 , respectivelyand could not be compared with the coefficients of other like-charged dissolved species. According to the general hypothesis of the SIT, 29 interactions between anionic complexes and Na + are often characterized by negative values of ε, while those for neutral complexes are considered as nil. The results in our previous paper questioned these classical assumptions. 8 Furthermore, research on this subject raised important issues that rely on restricted available references. The selected value 29 of ε(UO(CO3)2 2-,Na + ) = -(0.02 ± 0.09) kgw mol -1 was commonly taken as an analogy for ε(MUO2(CO3)3 2-,Na + ), and ε(M2UO2(CO3)3,NaX) was estimated to be nil, as postulated in SIT for non-charged species. 29,30 Comparison with other datai.e. ε(MgUO2(CO3)3 2-,Na + ) = (-3.0 ± 0.3) kgw mol -1 proposed by Dong and Brooks 10 and ε(Ca2UO2(CO3)3,NaClO4) = -2.7 kgw mol -1 evaluated from the work of Kalmykov and Choppin 2 -, produces equivocal implications of involved ion interaction processes, which led to the non-selection in Guillaumont, et al. 29 What also remains ambiguous is whether the affinity of Na + with CanUO2(CO3)3 (4-2n)- complexes is capable of yielding globally positively-charged complexes that may contribute to the relatively large and the positive values of ε in NaCl evidenced in Shang and Reiller, 8 as suggested in theoretical calculations. [31][32][33] For the sake of an exhaustive consideration, variations of the reaction constants should be measured in a larger range of ionic strength in which the potentially strong ion interaction processes could be better described. It is hoped that the completeness of reliable information with respect to thermodynamic constants in specific conditions would contribute to a deeper understanding of ion interaction mechanisms and fill the knowledge gap in the thermodynamic data.

Sodium perchlorate, when dissolved in water, can hardly react through coordination, or in any other manner, with different aqueous species. 34 The lack of chemical reactivity in solutions allows its extensive use as an inert background salt in the reactions of ionic materials. To the best of our knowledge, Kalmykov and Choppin 2 have used NaClO4 as the supporting electrolyte, and extended the ionic strength up to 3 m. The authors evidenced a noticeable deviation from the theoretical linear trend of log10 2.1.3 vs. the ionic strength when using SIT. 29 However, only Ca2UO2(CO3)3(aq) was postulated, and the monocalciumtriscarbonatouranyl complex was not identified. Furthermore, the lack of experimental points between 1 and 3 m NaClO4 ionic strength also adds difficulties to separate the electrolyte effects from the ligand effects, particularly important aspect of ion paring formation in the sequential binding processes between Ca 2+ and UO2(CO3)3 4-.

With increasing electrolyte concentration, the aqueous ions interactions are reinforced by the increasing mutually electrostatic effects disturbing the inner-and outer-sphere water molecules of solute and solvent ions and forming ion-pairs. 35 The cation-anion pairs can be divided into three categories according to the varying extents of ionic association: (i) the solvent-separated ion-pairs (SSIP) involve generally two layers of water molecules between the ion partners of the pair, and (ii) the solvent-shared ion pairs (SIP) contain one; (iii) the contact ion-pairs (CIP) are used to describe the situation where the desolvation of the inner-sphere water layer of both ions is inevitable. Several studies reported the hydration of Ca 2+ , Na + , ClO4 -and Cl -in aqueous solution and their tendencies to form different types of ion-pairs as a function of electrolyte concentration. It has been shown that the hydration layers of Na + and Ca 2+ can be depicted as Na + (H2O)4(H2O)9(H2O)5 and Ca 2+ (H2O)6-8(H2O)12, respectively. Na + is less hydrated in its first hydration sphere than Ca 2+ . The hydration of the larger anions ClO4 -and Cl -arereported to be 3.8 H2O and 6 H2O, respectively. [36][37][38][39][40][41] More specifically, Tu, et al. 42 attributed the broadening bands of ClO4 -in Raman spectra when increasing the Ca(ClO4)2 concentration to the growing number of SIPs. They also assumed that the oppositely charged ions are more likely to form CIPs if they are comparatively hydrated, suggesting the easier association between Ca 2+ and Cl - than between Ca 2+ and ClO4 -. The fundamental properties and chemical behaviour of the ions from the electrolyte would help us to clarify the ion interactions mechanism at high electrolyte concentrations and gain an in-depth explanation for the experimental results in NaCl and NaClO4.

The aim of this work is to determine the thermodynamic formation constants of CanUO2(CO3)3 (4-2n)-complexes in NaClO4 media using TRLS. Despite the intricate approaches, it seems that the effects of electrolyte ions may have no noticeable impact on the luminescence results over a certain range of ionic strength. The upper limit of NaClO4 ionic strength will be evaluated at the concentration above which the analysis of uranyl luminescence for speciation is no more valid. By analysing the evolution of formation constants within the Im range, this study is expected to clarify the CanUO2(CO3)3 (4-2n)--ClO4 -interaction mechanisms. These results will be discussed considering former observations in NaCl medium. 8 Variations of specific ion interaction coefficients as a function of electrolyte concentration will be described and predicted by different models proposed in the literature. Extension of the SIT approach would also be feasible to much higher ionic strength based on the present results.

Moreover, as regards to photophysical and radiochemical processes that possibly occurred when exciting UO2 2+ in perchloric environment, the most debated problem prevails on the reactivity of UO2 2+ in different statesi.e. excited or ground state. [43][44][45][46][47][48][49][50][51][52][53] It was proposed that the influence of perchlorate and hydrogen ions, especially the effect of quenching, was occurring at high ClO4 -concentrations. [51][52][53] However, what is less clear is the degree to which the luminescence mechanisms proposed in these studies can be transposable to the Ca-UO2-CO3 systems, as well as the sensibility of the overall packing scheme in the UO2-CO3 structure towards perchlorate interaction. The results in this work may raise intriguing questions about the quenching effect of perchlorate ions on triscarbonatouranyl units.

III -3 Experimental Section

Materials

An appropriate quantity of high purity U3O8 solid was dissolved in 60-70% perchloric acid (Sigma-Aldrich, ACS reagent, ≥ 99%) in Millipore deionized water (Alpha-Q, 18.2 MΩ cm -1 ).

All measurements were performed at the same uranium concentration of 50 μm.

Sodium hydroxide (Analytical Grade, Fisher) and diluted perchloric acid were used to adjust the aqueous pH values. The introduction of carbonate was performed with freshly prepared sodium bicarbonate solution (NaHCO3, Analytical Grade, Fisher). As all the sample solutions were open to the atmosphere during the preparation, carbon dioxide in the air equilibrated with CO2(aq), hydrogen and bicarbonate in solution. Therefore, the amounts of added NaHCO3 were calculated based on the equilibrium between the pH-fixed solution and atmospheric CO2(g) in the absence of calcium. In order to maintain the electrolyte concentrations at constant values, the volumes of reagents were kept small compared to that of the solution. Calcium perchlorate (Ca(ClO4)2:4H2O, Sigma-Aldrich, ACS reagent, ≥ 99%) solutions of various concentrations were prepared to control the concentration of calcium for each sample. All reagent preparations were conducted by weighing. Concentrations of involved species are expressed in molality units (mol kgw -1 or m) with squared-brackets in this work. Necessary prerequisite for further measurement-no variation of Ca concentrations has been verified at the present U(VI) concentration (data not shown). For every sample the saturation indices of UO3:2H2O(cr) and calcite are negative, which means that precipitation is unlikely.

Modification of pH electrode and measurements of liquid junction potential

A hydrogen ion-sensitive glass electrode (type, Mettler Toledo, USA) was used to perform the measurements of pH values. Considering the particular electrolyte anion in this work, the saturated KCl solution in the reference electrode was replaced with 3.5 m NaClO4 (2.99 M NaClO4 + 0.01 M NaCl), to avoid the KClO4 precipitation inside the liquid junction. The electrode was carefully handled in this procedure and then properly conditioned for reuse. The change of the reference inert solution generates a systematic deviation of potential difference between the electrodes. The potential gradient between the silver chloride electrode and the buffer solution in the pH electrode would remain constant, whereas that between the reference electrode and the test sample, separated by a semi-permeable membrane, needs to be re-determined for each electrolyte concentration. For this reason, a series of NaClO4 solutions of from 0.01 to 3.5 m acidified by HCl to pcH 2 (= -log10[H + ]) were prepared to measure the junction potentials with the reconditioned electrode of 3.5 m NaClO4 reference solution. 54 The potential values Emes read on pH-meter for the NaClO4 solutions of the same acidity (-log10[H + ] = 2) but different ionic strengths are reported in Table III -S1 of the Supplementary Information (SI). For a given inert electrolyte concentration, the difference between the value of Emes measured at Im = 0.01M and that at other Im is including the liquid junction potential of the reconditioned electrode and the activity coefficient of H + . Therefore, the experimentally measured ∆Emes were taken to correct the pH value of sample solutions at specific electrolyte concentrations.

A four-point calibration using the commercial buffer standard solutions of known and fixed pH values, respectively of 1.68, 4.01, 6.87 and 9.18 (SI Analytics, Mainz, Germany) was daily required to maintain the accuracy. 55 The method of ordinary least squares was employed to estimate the 95% confidence limits, corresponding to 0.05 pH units of maximum uncertainties. 56 Measurements of pH values were performed in triplicate at 22°C. For each test sample, the pH value was regularly checked and adjusted if necessary. The first pH measurement was carried out after adding appropriate quantities of NaHCO3(aq). The samples were then let open to air. After two hours, the pH value was checked for the second time and slightly adjusted if necessary. The pH never evolved after the nest additions of calcium perchlorate into the sample solution. The pH was then measured prior to each spectrum acquisition.

Time-resolved laser-induced luminescence spectroscopy

A 1 cm quartz cuvette (QS101, Suprasil, Hellma Analytics) containing a 2 mL aliquot of U(VI) test sample was used to measure the luminescence spectra in the wavelength range from 415 to 615 nm. The fluorescence cuvette was placed in the thermostated sample holder, connected to the thermostat (Peter Huber Kältemaschinenbau AG, Germany) with circulating water, during at least 10 min before data acquisition. The cooling water temperature was set at 22ºC for minimizing the temperature fluctuation on uranyl fluorescence. A frequency-tripled Nd:YAG laser at 355 nm (Surelite, Continuum, USA) working with 5 ns pulse at 10 Hz of about 170 mJ as energy was used as the pump laser of an optical parametric oscillator system (Horizon, Continuum, USA). The excitation wavelength was set at 450 nm for all luminescence measurements. This wavelength represents the better compromise between the selective excitation of uranium(VI) and the excitation energy delivered by the optical parametric oscillator. A RJP-734 Joule-meter (Laser Probe, Inc.) was used to monitor the excitation energy.

The 300 lines mm -1 grating, favouring the convolution of various fluorescing components with the detection apparatus, was selected in the monochromator spectrometer (Acton 300i, Roper Scientific, U.S.A.). The resulting measured fluorescence signal was detected by an intensified CCD camera (Andor, U.K.) maintained at -10°C by a Peltier effect. All emission spectra acquisitions started at the delay time of 25 ns with the gate width fixed at 1 μs.

The present concentration of U(VI) was low enough not to generate any pre-and post-filter effects. The observed fluorescent behaviour was thus assumed to be the cumulative contribution of aqueous U(VI) complexes, on the condition that neither precipitation nor colloidal aggregation occurred in solution. The recorded spectra were integrated over the emission wavelength span to evaluate the fluorescence intensity at various delay times. The luminescence intensity at D = 0, as well as luminescence decay-time, were derived by fitting the intensity results to the exponential fluorescence decay function, expressed as follows, [57][58][59][60] FI = ∑ FI 0 τ i exp (-

D τ i ) i (1 -exp (- W τ i )) (III -1)
where FI and FI0 are respectively the fluorescence intensity at a specific delay time and at D = 0. The gate width W equals to 1 μs and τi represents the decay-time of species i.

Origin 9.0 software (OriginLab Corporation, USA) was used to analyse experimental data.

The background noise was recorded and subtracted from measured fluorescence emission spectra for each series prior to subsequent peak analysis. The trapezoidal integration method was used to calculate the fluorescence intensity. The least-square fitting process was then evaluated by the Levenberg-Marquardt algorithm.

Correction for slope analysis and calculation of related Ringböm coefficient.

For explanatory purposes, the equations used for the slope analysis and the meanings of the corrective factor α, also the Ringböm 61 coefficient relative to UO2(CO3)3 4-, are recalled here.

The calculation procedure for α has been exemplified in our previous study. 8 The Ringböm coefficientcalculated using the thermodynamic constants from Table III

K n.1.3 = [Ca n UO 2 (CO 3 ) 3 (4-2n)-] [Ca 2+ ] n [UO 2 (CO 3 ) 3 4-] n = 1, 2 (III -2)
where the concentrations are expressed in molality (m).

illustration, fluorescence decay spectra of UO2-CO3 sample at pH 7.6 were shown in the Fig.

III -S2 in our previous study. 8 The spectral evidence indicated that FI0 (UO2(CO3)3 4-) cannot be directly measured at lower pH values (pH < 8). The Ringböm coefficient α was therefore introduced to deduce FI0 (UO2(CO3)3 4-) at desirable pH values (7 < pH ≤ 9), because there is no difficulty to acquire FI0 (UO2(CO3)3 4-) at pH 9 at the investigated ionic strengths (0.1 m ≤ Im ≤ 3.5 m). We expressed the slope analysis as in eqn 3), R is the concentration ration between the [CanUO2(CO3)3 (4-2n)-] and [UO2(CO3)3 4-] with the factor α to evaluate the amount of UO2(CO3)3 4-at the pH values different from pH 9, at which the FI0(UO2(CO3)3 4-) was measured.

The slope n represents the stoichiometric number of calcium ion and the intercepts log10 Kn. 1.3 stand for the formation constants of CanUO2(CO3)3 (4-2n)-.

log 10 R = log 10 [Ca n UO 2 (CO 3 ) 3 (4-2n)-] [UO 2 (CO 3 ) 3 4-] α ⁄ log 10 R = log 10 FI 0 (Ca n UO 2 (CO 3 ) 3 (4-2n)-) FI 0 (UO 2 (CO 3 ) 3 4-) α ⁄ = = log 10 K n.1.3 + n log 10 [Ca 2+ ]
(III -3)

Specific ion interaction Theory (SIT)

The SIT approach is used to extrapolate the conditional stability constants to infinite dilution, as the long-range and short-range forces are both considered in this model. 29 According to the specific ion interaction theory, the expression for the activity coefficient is in the form of Eq. III -4:

log 10 γ j = -z j 2 D H + ε(j, k, I m ) m k (III -4)
with DH the Debye-Hückel term,

D H = A√I m 1+1.5√I m (III -5)
where ε(j,k) (kgw mol -1 ) represents the ion interaction coefficient between the aqueous ion j and either the cation or the anion k of background electrolyte (depending on the whether the ion j is cationic or anionic) of concentration Im on the molal scale.

The value of 0.507 kg 1/2 mol -1/2 was chosen for Debye-Hückel constant A, instead of 0.509 kg 1/2 mol -1/2 calculated for 25°C, because the measurements were conducted at 22°C. 29 The relationship between conditional formation constants and ionic strengths is expressed by Eq. III -6. where

nCa 2+ + UO 2 2+ + 3CO 3 2-⇄ Ca n UO 2 (CO 3 ) 3 (4-2n)- β n.1.3 = [Ca n UO 2 (CO 3 ) 3 (4-2n)-] [Ca 2+ ] n [UO 2 2+ ][CO 3 2-] 3 n = 1,
Δz 2 = z(Ca n UO 2 (CO 3 ) 3 (4-2n)-) 2 -nz(Ca 2+ ) 2 -z(UO 2 2+ ) 2 -3z(CO 3 2-) 2 (III -7)
(Δz 2 = -16 for n = 1 and Δz 2 = -24 for n = 2) and

Δε = ε(Ca n UO 2 (CO 3 ) 3 (4-2n)-, Na + ) -nε(Ca 2+ , ClO 4 -)-ε(UO 2 2+ , ClO 4 -) -3ε(Na + , CO 3 2-)(III - 8)
Table III -2 summarizes the specific ion interaction coefficients used in calculations. It should be noted that in Guillaumont, et al., 29 the value of ε(UO2(CO3)3 4-,Na + ) = (0.08 ± 0.11) kgw mol -1 was assigned using the interaction coefficient ε(UO2 2+ ,ClO4 -) = 0.46 kgw mol -1 , and not from ε(UO2 2+ ,Cl -) = 0.21 kgw mol -1 , since the complexation of chloride towards dioxouranium(VI), though very weak, should be taken into consideration. In our previous study of NaCl medium, the use of ε(UO2 2+ ,Cl -) = 0.21 kgw mol -1 was avoided by applying the linear regression to the variation of log10 Kn.1.3 vs. Im, as an alternative to log10 n.1.3. 8 Therefore, the thermodynamic constants and ion interaction coefficients in these two works are both deduced with ε(UO2 2+ ,ClO4 -) = 0.46 kgw mol -1 .

III -4 Results and discussion

Evaluation of equilibrium constants and specific ion interaction coefficients

The upper limit of investigated ionic strength is enlarged to 3.5 m NaClO4, also in the Im range of application of SIT. Seven series of solutions are separately prepared at 0. point, are too small to be presented, due to the use of trapezoidal rule in the numerical integration method to deduce FI0 (the fluorescence intensity at D = 0). Weighted least-squares regression analysis was used to perform all linear fits and subsequently to determine the related parameters, i.e. slope and intercept mean values with 95% uncertainties. On account that the analytical calcium concentration is much larger than that of uranium, the aqueous Ca 2+ is considered as free calcium ion in linear plots. According to the meanings of parameters in linear regression analysis, the slope values, ca.1 and 2, correspond to the ratio of coordinated calcium towards the triscarbonatouranyl moiety. After rounding off the slope values to the integers of 1 and 2, the associated intercepts represent the formation constants of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) at a specific ionic strength. It is noteworthy that the fixed slopes intrinsically influence the uncertainties of associated intercepts. In the further analysis, the deduced uncertainties after imposing the slope values to integers would be adopted for the aim of consistency. -S1 in the SI). At the highest ionic strength, the addition of Ca 2+ does not give rise to the enhanced fluorescence intensity, or the second slope. For the sake of conformity, the results of Im = 3.50 m NaClO4 would not be included in the linear regression for evaluating the global stability constants and specific ion interaction coefficientsvide post for discussion. According to the experimental results listed in Table III -3, one can conclude that higher concentrated electrolytes require a larger quantity of Ca 2+ to the formation of ternary species.

This fact results from the increased electrostatic attraction between the salt ions and oppositely charged particles at higher ionic strengths. The conditional formation constants derived from the rounded off slopes in linear regression analysis are given in Table III Specifically, the values of Δlog10 K at 0.1, 0.5, and 1 m NaClO4 can be compared with the corresponding ones in NaCl medium. In the most dilute supporting electrolyte (0.1 m), there is nearly no difference in results, Δlog10 K 0.1 m = 2. From the slope in Fig. III -3 (a,b) and ε values in Table III -2, the specific ion interaction coefficients are estimated to be ε(CaUO2(CO3)3 2-,Na + ) = (0.02 ± 0.04) kgw mol -1 and ε(Ca2UO2(CO3)3,NaClO4) = (0.18 ± 0.07) kgw mol -1 , which are somewhat different from the NaCl case. 8 This will be discussed afterwards.

Comparison and interpretation of log10 ° and ε determined in NaCl and NaClO4

The comparison of log10 1. The deduced values of ε(CaUO2(CO3)3 2-,Na + ) and ε(Ca2UO2(CO3)3(aq),NaX) -X being either ClO4 -or Cl -in all hypothesisare summarized in Table III NaClO4. The ion interaction parameters of the ternary complexes were determined from these values and the change of aqueous medium anion from Cl -to ClO4 -leads to a noticeable decrease in the values of ε for the same complex.

The variations of specific ion interaction coefficients do show some interesting patterns that are tightly related to ionic strength and inert medium. The discussion of inconsistencies in ε values for the same counter-ion in different background electrolyte can conveniently starts by examining the appropriateness of the SIT model to explain the observed evolution. The Debye-Hückel theory is basically a theory of weak interactions, considering the long-range interaction in dilute binary solutions between homogeneously charged spherical ions. 63 The specific ion-interaction contribution arises from an Im-dependent second virial coefficient-type expression, originally proposed to account for the interactions between pairs of molecules for gases. 32,64 Nevertheless, ion association caused by ion-solvent and short-range interactions should not be ignored at high electrolyte concentrations. The enhanced formation of ion-pairs between Ca 2+ and ClO4 -or Cl -could enhance the competition with the interaction between Ca 2+ and UO2(CO3)3 4-. In this respect, the binding processes of Ca 2+ with structurally asymmetric UO2(CO3)3 4-group, calculated with theoretical methods, 65,66 are more complicated at higher Im, because the less evenly distributed charges can distort and polarize the immediate electron cloud of UO2(CO3)3 4-that confronts the assumption of SIT.

The quantitative comparison of ∆ε in different media draws our attention to the relevant factors that influence the activity of calcium during coordination. In our studied systems, the existing binary interactions implicate Ca 2+ -UO2(CO3)3 4-, Na + -UO2(CO3)3 4-, Ca 2+ -CaUO2(CO3)3 2-, Na + -CaUO2(CO3)3 2-, and Ca 2+ -ClO4 -/Cl -. Questions are raised about to which extent Na + is competitive to bind with UO2(CO3)3 4-and CaUO2(CO3)3 2-compared with Ca 2+ , and whether Ca 2+ is encumbered by the ion-pair formation with large anions. Li, et al. 31 examined the role of Na + on solvation of UO2(CO3)3 4-and CanUO2(CO3)3 (4-2n)-by classical molecular dynamics simulations. They reported that the presence of Na + is capable of alternating the coordination number of uranyl moiety and probably resulting in changes of global charge for UO2(CO3)3 4-and CaUO2(CO3)3 2-, but not likely to provoke a substitution from CaUO2(CO3)3 2-to NaUO2(CO3)3 3-. They provided indication for the Na + influence on inner-sphere level of uranyl unit and shed light on its key role of common ions in formation processes of CanUO2(CO3)3 (4-2n)-.

Regarding the ion association of Ca 2+ -ClO4 -/Cl -, SIPs are possibly the most predominant for Ca 2+ -ClO4 -at the investigated ionic strengths, as the inner hydration shell of Ca 2+ was observed intact up to 6 m in the work of Rudolph and Irmer. 67 These authors also concluded that for Ca 2+ -Cl -, the ion association is slightly stronger than that of Ca 2+ -ClO4 -, which can contribute to an increasing amount of CIPs. These interactions cannot be discounted in treating experimental data obtained in concentrated solutions. The differences in ∆ε in NaCl and NaClO4 may partly be explained by the ion triplet interactions of Na + -Ca 2+ -ClO4 -/Cl -that could reduce the activity of Ca 2+ to different degrees depending on the kind of anion with increasing ionic strengths. It would be instructive to consider the Pitzer 64 model to calculate the activity coefficient for this intractable system. Unfortunately, the fitting parameters for the ion partners of the pairs, especially for the UO2(CO3)3 4-species, are not completely available. Even if estimations of  (0) ,  (1) , and C Φ have been proposed from SIT parameters, 68,69 the  (2) coefficient in Pitzer expression used to describe the ion pairing effects was questioned for its degree of sophistication. 70 Therefore, the effort to extend this method to higher Im and acquire additional experimental data could be helpful. It would be interesting to perform measurements in a liquid medium supposed to be less complexing than ClO4 -and exhibit weaker quenching effect than Cl -e.g., sodium triflate. 71,72 Another aspect concerning the affinity of Na + with CanUO2(CO3)3 (4-2n)-can also make a distinction between ∆ε in NaCl and NaClO4. In previous theoretical calculations, the close-by Na + was suggested to indirectly interact with CanUO2(CO3)3 (4-2n)-, which may result in alternation of global apparent charge of the ternary complexes. The complex-forming tendencies could give rise to the formation of Na[Ca2UO2(CO3)3] + under high Im conditions. 33 Ciavatta 73 stressed the role of water molecules rearrangements in solvation sphere during complexation for interpreting chemical behaviour of metal ions at high ionic strengths. By comparing the interaction coefficients of a series of metal ions in several common salts, it was stated that the values of ε between oppositely charged ions usually have lower values if the ion partners are of similar size, because of the more easily occurred desolvation in both hydration shells, and further facilitated association between them. Considering its larger size, ClO4 -should be more likely to interact than Cl -with a potentially large-sized Na[Ca2UO2(CO3)3] + species.

The comparison of ε values well supports the hypothesis that the interaction between the CanUO2(CO3)3 (4-2n)-and background electrolyte ions is likely impacted by the integration of Na + implying that the neutral or positively charged quaternary complexes would more likely interact with the increasing concentrations of medium anions rather than with alkali cations. Otherwise, the interaction coefficients determined respectively in NaCl and NaClO4 solutions should not significantly differ from each other. Nevertheless, there is a need of more direct evidences, e.g.

based on charge determinations.

Evolution of spectral-and temporal-luminescence parameters

The evolutions of emission spectra of aqueous Ca-UO2-CO3 ternary species are illustrated in Fig. III -1. The spectra were decomposed using a Gaussian distribution function. As already observed for the NaCl electrolyte, no apparent shift in peak positions was observed in NaClO4 using the 300 lines mm -1 grid. Moreover, the relatively high ionic strength in the case of Im = 3.5 m NaClO4 did not induce any shift in wavelength. Peak maxima were determined at 464.4 ± 0.7 nm, 483.6 ± 0.7 nm, 503.8 ± 1.1 nm, 525.4 ± 1.9 nm, and 547.9 ± 4.9 nm (2σ) by averaging all peak maxima of 114 sample spectra, and agree well with literature values. 2,3,6,8 The luminescence of the triscarbonatouranyl corresponds to transitions from non-bonding 5f orbitals of uranium to a mixed orbital with uranyl-σ (5f, 6d) and oxygen-2p character. 74,75 The energy levels of the excited and ground states are almost insensitiveat least at the resolution of the apparatus using the 300 lines mm -1 grid, which gives a difference of 0.2 nm between each pixel of the camerato the presence of calcium located out of the first coordination sphere, composed of three bidentate carbonate ligands. 15 In order to probe the role of long-range solvent anions and examine the effects of electrolyte concentration, a comparison of several normalized emission spectra is shown in Fig. III -4. The species percentages of sample solution were calculated using the log10° and ε determined in this work for NaClO4 medium and the results determined in our earlier work for NaCl medium. 8 Four solutions that contain nearly identical composition in binary/ternary species are selected complexes. Table III -6 lists the percentages of major U(VI) complexes in selected solutions. Subtle spectral differences have been observed when comparing these spectra. The dicalcium(II) complex contributes to better resolved spectra and narrower peak widths than the monocalcium(II)triscarbonatouranyl(VI) complex, as observed elsewhere, 3,6 and also in our previous work in NaCl medium. 8 In particular, the ratio of maximum heights between the second (ca. 484 nm) and the third peak (ca. 504 nm) varies as a function of complexed calcium concentration, as illustrated in Fig. III -6. These two peaks represent the luminescence transitions from the lowest triplet excited state (10) to different vibrational levels of singlet ground state (0υ), with υ = 0 for the peak ca. 484 nm and υ = 1 for ca. 504 nm. 76,77 The decreasing trend of ratio values, though relatively weak, suggests that the 483.6 nm peak (10 → 00 at 20,680 cm -1 ) is more sensitive to the calcium complexation than the 503.8 nm (10 → 01 at 19,850 cm -1 ) peak. The vibronic frequency corresponding to the symmetrical U-O stretching mode can be determined from the difference in energy between the 10 → 0υ (υ = 0-3, 483.6 nm, 503.8 nm, 525.4 nm, 547.9 nm) transitions, and was found to be about 810 ± 19 cm -1 which is in good accordance with the symmetrical stretching υs values measured by Raman/IR spectroscopies. 78 The decay-times of CaUO2(CO3)3 Regarding the question of quenching effect of perchlorate ions, the research to date has mostly been focused on the spectral and temporal behaviours of the aqueous uranyl ions in HClO4/NaClO4 environment. 51,52,71 It was reported that the decay-time of UO2 2+ depends significantly on the acidity and perchlorate concentration. The increase of [ClO4 -] would lead to a decrease of decay-time at constant pH value, while the increase of acidity would contribute to longer decay-times of UO2 2+ except at very high concentration of HClO4 (> 10 M). 71 In the work of Rustenholz et al. 52 the concomitantly decreasing trends of fluorescence intensity and decay-time were observed for the uranyl ion when increasing perchlorate concentration at the same [H + ]. Interestingly, the decreases of fluorescence parameters FI and τ, both exhibit an obvious steep at [ClO4 -] = 2.75 M, i.e. 3.14 m. 79 The discussion on these results typically examines the quality of fit by chemical-or physical-based models. However, the conclusions in literature also could seem arguable to apply directly to the Ca-UO2-CO3 system since the published research mainly focuses on free uranyl ions in an extremely acid environment in which the perchlorate ion approaches closer to the UO2 2+ central entity. The present luminescence results for CanUO2(CO3)3 (4-2n)-(n = {0; 1; 2}) does not sit on a sufficiently wide basis of evidence to justify the impact of aqueous acidity. The quenching effect of ClO4 -still remains as a plausible reasoning for interpreting the different variations of FI and τ observed at Im = 3.50 m NaClO4, even though its influence on U(VI) luminescence behaviour might be very subtle due to the protection of three carbonates around the uranyl moiety.

As shown in Fig. III -8, the invariance of UO2(CO3)3 4-decay-time ca. 10 ns vs. Im provides evidence for its insensitivity of deexcitation processes to the [ClO4 -]. For the ternary species, the decay-time of Ca2UO2(CO3)3(aq) manifests higher increasing tendency than CaUO2(CO3)3 2- at Im = 2.46 m NaClO4. We could infer that under mildly alkaline environment (pH = 9), luminescence properties of the large negative ion UO2(CO3)3 4-, caged by Na + in its surrounding environment and undergoing repelling forces with ClO4 -, is less likely to be affected by ClO4 -.

In the presence of calcium, the earlier mentioned hypothesis that the affinity of Na + is capable of altering the global apparent charge seems to reconcile with our experimental results and should be reconsidered in the light of molecular dynamics modelling. 31,33 A case in point is proposed formation of the positively like-charged Na[Ca2UO2(CO3)3] + . If occurring, the effective distance between aqueous U(VI) complexes and ClO4 -would be shortened due to the attractive force between them. This could give the reason for the experimental observations that the quenching effect of ClO4 -started to be noticeable at high Im.

As illustrated in Fig. III -S1 of the SI, the suppression of fluorescence intensity is more evident than decay-time on addition of Ca(ClO4)2 into the system of Im= 3.5 m to the extent that the intensity levelled off and the second slope no longer appeared, but meanwhile the decay-times showed a continuous increase. It is thus suggested that the quenching effect of ClO4 -is more effective in diminishing fluorescence intensity than decay-time.

Description of equilibrium constants variation by different models

Using our newly determined constants and parameters (log10 ° and ε), the predominance diagrams at studied ionic strengths are constructed using PhreePlot 80 which is including the program PhreeqC 81,82 that contains the ThermoChimie database, 83 This differences between simulations and experimental results may suggest the possible variation of ε as a function of ionic strength, since the specific ion interaction coefficients, implemented in the ThermoChimie database, are taken as invariable up to the upper limit of validity of the SIT approach, ca. Im = 3-4 m, in the PhreeqC calculations. The back extrapolation from infinite dilution using log10 °n.1.

ε(j, k, I m ) m k = εI m -log 10 (1 + K I m ) (III -9) 𝜀(𝑗, 𝑘, 𝐼 𝑚 ) = 𝜀 ∞ + 𝜀 0 -𝜀 ∞ 𝐼 𝑚 +1
(III -10)

ε(j, k, I m ) = ε 1 + ε 2 log 10 I m (III -11)
where Eq. III -9 and Eq. III -10 are reported by Bretti and co-workers, 84,85 and Eq. III -11 is proposed by Ciavatta, 73 and also reported in e.g. Guillaumont et al. 29 and used by Brown and Ekberg. 86 is reasonably lower than that of one-parameter SIT model. The convergence of model fitting is finally attained even with the limitation of number of experimental points. It is therefore suggested that the new parameters in two-parameter models are viable to generate trend lines for current data to higher ionic strengths. With the mainly illustrative purpose, the range of electrolyte concentration is extended to 6 m, at which the use of two-parameter model has been proven compatible with the traditional one-parameter SIT model for classical 1:1 type electrolytes. 85 One can observe that no important inconsistencies are identified over the ionic strength from 0.1 to 2.46 m. The general trend can be well fitted by the one-parameter SIT model over this range, but the quality of the fit is degraded compared to the two-parameter models. In a larger range of ionic strength up to 6 m, the difference between the two models proposed by Bretti, et al., 84,85 is almost negligible. In their papers, this approximation was also verified and considered acceptable within a certain electrolyte concentration. However, the difference between the model proposed by Ciavatta 73 and the simplified SIT equation increases considerably for ionic strength above 3 m. The dependence of ion interaction coefficient on ionic strength would give rise to more significant differences in resulting conditional stability constants at higher electrolyte concentration. Concerning the residual sum-of-squares, generally used to quantify the goodness-of-fit, it is shown that the two-parameter models fit the same data set better than the linear model.

Typically, multi-parametric nonlinear regressions give more flexibility to the curve fitting and

III -7 Supporting Information

IV -1 Abstract

The formation constants and specific ion interaction coefficients of MgUO2(CO3)3 2-complex were determined in 0.1 to 1 mol kgw -1 NaCl and 0.1 to 2.21 mol kgw -1 NaClO4 media in the framework of the specific ion interaction theory (SIT), using time-resolved laser-induced luminescence spectroscopy. The upper limits of ionic strength were chosen in order to limit

IV -2 Introduction

The alkaline earth triscarbonatouranyl complexes MnUO2(CO3)3 (4-2n)-especially for Mg and Caseem to have an essential role in controlling the U(VI) speciation and mobility in mildly basic carbonate-rich groundwater by increasing U(VI) solubility and facilitating its migration. [1][2][3][4][5][6][7][8] Due to the importance of these species, their thermodynamic constants and functions -log10β°n.1.3, ∆rG°m ∆rH°m, ∆rS°m -, as well as their ionic strength dependence, have been investigated by different methods such as time-resolved luminescence spectroscopy (TRLS), 1,2,9-13 ion-exchange, 14,15 UV-Visible spectroscopy, 16,17 and potentiometry. 18 Even if the existence of the alkaline earth carbonate phases of uranium(VI) are known since the 19 th century, 19,20 the first evidence of the Ca aqueous complexes is quite recent in uranium chemical history, 1 see e.g. in the critical compilations of thermochemical data commissioned by the Nuclear Energy Agency from the Organization for Economic Co-operation and Development -NEA-OECD [21][22][23] . The aqueous complexes had not been evidenced yet in Grenthe, et al., 21 and were not selected in Guillaumont, et al. 22 A selection of thermodynamic functions and constants have been proposed recently in Grenthe, et al., 23 , based on determinations available before 2018the most recent reviewed work being Lee, et al. 10 The only works that were dealing with ionic strength variation are Kalmykov and Choppin 11 for Ca, and Dong and Brooks 15 for Mg, which have not been judged satisfying in Guillaumont, et al. 22 and Grenthe, et al., 23 respectively. The selected data, obtained in a limited range of ionic strength, could not allow proposing both extrapolations of formation constants to infinite dilution (log10β°) and the specific ion interaction coefficients (ε), using the specific ion interaction theory (SIT), which is the framework of the proposition of thermodynamic constants for the NEA-OECD commissioned reviews.

In recent works, we reported the conditional formation constants for CanUO2(CO3)3 (4-2n)-species at Im = 0.1 to 1 mol kgw -1 in NaCl, 24 and 0.1 to 2.21 mol kgw -1 in NaClO4, 25 which had led to the propositions of log10β°(CaUO2(CO3)3 2-), log10β°(Ca2UO2(CO3)3(aq)), ε(CaUO2(CO3)3 2-,Na + ), ε(Ca2UO2(CO3)3(aq),NaCl), and ε(Ca2UO2(CO3)3(aq),NaClO4). These works could not have been included in the latter NEA-OECD review, 23 as they were published after the peer-review.

Particularly, we showed that two usual hypotheses concerning the value of ε for these complexes were not justified: (i) the charge-like analogy of CaUO2(CO3)3 2-and UO2(CO3)2 2-was not justified concerning their ε value towards Na + ; and (ii) the nil value of ε for neutral complexes could not be generalized for these large speciesone can also notice that Ciavatta 26 already represented non-nil ε values for Cd(II) and Hg(II) neutral complexes.

The spectroluminescence titration of uranyl carbonate solutions with alkaline earth metal cations at varying pH valuesin order to avoid the precipitation of schoepite and calcitewas proven feasible and robust to derive the log10β° and ε values by extrapolating log10β measured at specific ionic strengths to infinite dilution using SIT. In NaCl, 24 the striking quenching effect of chloride impeded the acquisitions of well-resolved luminescence spectra at Im > 1 mol kgw -1 .

In NaClO4, 25 the more limited quenching effect of perchlorate still restricted the Im upper limit to 2.46 mol kgw -1 . The interaction coefficients of CanUO2(CO3)3 (4-2n)-complexes exhibited drastically different patterns depending on the nature of the media. The positive interaction coefficient values were estimated to be ε(Na + ,CaUO2(CO3)3 2-) = (0.29 ± 0.11) kgw mol -1 and ε(Ca2UO2(CO3)3(aq),NaCl) = (0.66 ± 0.12) kgw mol -1 . The change of background electrolyte to NaClO4 gave rise to less positive values as ε(Na + ,CaUO2(CO3)3 2-) = (0.02 ± 0.04) kgw mol -1 and ε(Ca2UO2(CO3)3(aq),NaClO4) = (0.18 ± 0.07) kgw mol -1the value of the neutral complex being significantly different from 0.

Our previous results provided evidence for the medium-dependent interaction coefficients of a specific complex. The obtained ε values for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) suggest that the hypothesis ε(CaUO2(CO3)3 2-,Na + ) = ε(UO2(CO3)2 2-,Na + ) = (-0.02 ± 0.09) kgw mol -1 requires particular attention, as well as the nil value for ε(Ca2UO2(CO3)3(aq), NX) -NX being the background electrolyte. In addition, the log10βn.1.3 vs. Im functions displayed remarkably different patterns in NaCl and NaClO4 that are possibly accounted for by the tertiary ion-ion interactions as well as ion pairing at high ionic strengths. The interactions between Ca 2+ and UO2(CO3)3 4-are suspected to be perturbed by the ion association of Ca 2+ with other competing anions in the background solutions. 27 Moreover, the affinity of Na + can induce a distortion of carbonate denticity from bidentate to monodentate and thus change the first shell coordination number of U, as proposed by molecular dynamics simulations. 28,29 These findings for Ca-UO2-CO3 species strongly suggest the important influence of high ionic strengths on Ca 2+ /UO2 2+ /CO3 2-complexation.

In this work, we propose new determinations of thermodynamic constants of formation for another environmentally relevant alkaline earth cation, Mg 2+ , to form the MgnUO2(CO3)3 (4-2n)-complexes at ionic strength 0.1 to 1 mol kgw -1 NaCl and 0.1 to 2.21 mol kgw -1 NaClO4 media, using slope analysis of the uranium(VI) spectra in time-resolved laser-induced luminescence spectroscopy. The conditional constants will be extrapolated to infinite dilution in the framework of SIT providing both log10β° and ε values for the evidenced complexes.

Even if Mg 2+ and Ca 2+ are likely to have a similar role in uranyl carbonate complexes, different hydration structures of these two elements still lead to various degrees of difficulty in complexation with the UO2(CO3)3 4-moiety. 18,[30][31][32][33][34] Especially the more difficult water exchange between Mg(H2O)6 2+stable octahedral with perpendicular square plan 31,33,34 and UO2(CO3)3 4-. Jo, et al. 16,17 evidenced that the formation of Mg-UO2-CO3 complexes is subjected to endothermic reaction, whilst that of Ca-UO2-CO3 complexes is exothermic, which was also observed by Endrizzi and Rao. 18 Furthermore, as for other ligandssee Fig The thermodynamic functions of MgnUO2(CO3)3 (4-2n)-and CanUO2(CO3)3 (4-2n)-could be important to account for when describing the geochemistry of representative cases where the concentration of magnesium is at least comparable or higher than that of calcium 7 . Using the estimation from Dong and Brooks 14 for MgUO2(CO3)3 2-and CaUO2(CO3)3 2-the equality of concentration of the two complexes should occur at [Mg 2+ ]/[Ca 2+ ] ca. 12, and ca. 24 using the Dong and Brooks 15 constants extrapolated using Davies 35 equation for Mg.

Reiller and Descostes 7 identified some cases where the need of dedicated ionic strength determinations for both Ca and Mg complexes were stressed: the clay-rich minerals dedicated to geological repository for radioactive wastese.g. the Boom Clay in the Belgian context, 36 and the Callovo-Oxfordian Clay in the French context 37 -, and seawater. 38 We can select other instances as: a Canadian aquitard 39,40 characteristic of the spatially varying ratio Mg/Ca with depth; and in addition, the cases of uranium-rich Indian groundwaters 41,42 , sources of local drinking water, as a part of discussion considering its humanistic and ecological significance.

Theoretical modelling using PhreeqC 43,44 are performed by implementing the newly measured log10β° and ε and those for CanUO2(CO3)3 (4-2n)-in our previous studies 24,25 . Clear differences in speciation results are analysed after comparing the modelling results with other available thermodynamic constants in literature.

IV -3 Experimental Section

Materials

Stock solutions of U(VI) were separately prepared by dissolving appropriate quantities of U3O8 solid in 60-70% perchlorate acid (Sigma-Aldrich, ACS reagent) for two days, and in 37% hydrochloric acid (Sigma-Aldrich, ACS reagent) for one week. The concentrations of U(VI)

were measured by inductively coupled plasma mass spectroscopy (ICP-MS). The final concentration of uranium was diluted to 50 µmol kgw -1 in both ionic media for further experiments. Anhydrous NaClO4 (Sigma-Aldrich, ACS reagent, ≥ 99%) and NaCl (Sigma-Aldrich, ACS reagent, ≥ 99%) were used to prepare the inert electrolytes with Millipore deionized water (Alpha-Q, 18.2 MΩ cm -1 ) at varying ionic strengths, respectively 0.1, 0.51, 1.05, 1.62 and 2.21 mol kgw -1 for NaClO4 series, and 0.1, 0.25, 0.50, 0.75 and 1.0 mol kgw -1 for NaCl series.

Freshly prepared NaHCO3 (Analytical Grade, Fisher) solution was used to introduce carbonate into the aqueous solutions prior to the addition of alkaline earth metal ions. The amounts of added bicarbonate solution were previously calculated considering the equilibrium between the atmospheric carbon dioxide and the planning pH values. Magnesium perchlorate (Mg(ClO4)2, Sigma-Aldrich, ACS reagent, ≥ 99%) and magnesium chloride (MgCl2, Sigma-Aldrich, ACS reagent, ≥ 99%) were used to control the concentration of magnesium. The addition of reagents was restricted to be in small volume so as not to disturb the electrolyte concentrations. The reagent preparations were performed by weighing. The species concentrations were expressed in molality units (mol kgw -1 or m in the rest of the text). The saturation indices of UO3:2H2O(cr) and MgCO3 from the ThermoChimie 10a databasesee Grivé, et al. 45 ; http://www.thermochimie-tdb.comfor every sample were checked as negative or near zero values, which suggested undersaturation of schoepite or magnesite during experiments.

The adjustment of aqueous pH value was conducted with NaOH (Analytical Grade, Fisher) and diluted HCl or HClO4 solutions. Two hydrogen ion-sensitive glass electrodes (Mettler Toledo, USA) were used for pH measurements. As for the Ca-UO2-CO3 study, 25 the pHelectrode used in NaClO4 was reconditioned by changing the filling solution from saturated KCl to 3.5 m NaClO4 in order to avoid the precipitation of KClO4 inside the frit. The junction potential for these two electrodes were estimated by measuring the potential values of specially prepared NaCl and NaClO4 solutions of the same log10[H + ] but different Im. The establishment of linear relationship between the measured potential values and the ionic strengths is important to correct and convert the measured potential to actual pH value for each sample solution. 46 The experimentally measured ∆Emes (difference between Emes measured at Im = 0.01 m and that at other Im) were reported in Table IV -S1 of the SI for NaCl medium. The results for NaClO4 are available in Table IV -S1 of the SI in Shang, et al. 25 Commercial solutions of pH = 1.68, 4.01, 6.87 and 9.18 (SI Analytics, Mainz, Germany) were daily used to check the calibration and maintain the accuracy. A 95% confidence interval that corresponds to 0.05 pH units of uncertainties was obtained by the method of ordinary least squares. All pH measurements were performed in triplicate at 22°C.

Time-resolved laser-induced luminescence spectroscopy

The time-resolved luminescence apparatus was already described in details elsewhere. 24,25 The luminescence spectra were acquired at the excitation wavelength of 450 nm, assuring a selective excitation of U(VI), 47 using a frequency-tripled Nd:YAG laser at 355 nm (Surelite, Continuum, USA) as the pulsed laser beam and an optical parametric oscillator (OPO) system (Horizon, Continuum, USA). The excitation energy was monitored by a RTP-734 Joule-meter (Laser Probe, Inc.). The temperature of sample solutions was set at 22°C during all the measurements using the thermostated sample holder, connected to a thermostat (Peter Huber Kältemaschinenbau AG, Germany) with circulating cooling water to minimise the temperature influence on the U(VI) luminescence. 48 The luminescence signal was collected by an Andor monochromator spectrometer using either the 300 or the 1800 lines mm -1 grating. The U(VI) spectra were recorded in the emission wavelength range 415-615 nm using an intensified CCD camera (Acton).

The fluorescence behaviour was considered to be cumulative contribution of different U(VI)

species. The time-resolved emission spectra acquisitions were acquired at different delay time D with the gate-width W fixed at 1 μs, and then analysed using Origin 9.0 software (OriginLab Corporation, USA), including subtraction of background noise, and peak integration with the trapezoidal rule. The intensity results were fitted to the exponential decay function to derive the luminescence intensity at D = 0 and the luminescence decay-time, expressed as already discussed, 24,25 FI = ∑ FI 0 τ i exp (-

D τ i ) i (1 -exp (- W τ i )) (IV -1)

IV -4 Results and discussion

Evaluation of equilibrium constants and specific ion interaction coefficients

A series of solutions was prepared at desired ionic strengths with the maximum electrolyte concentrations fixed to 1 mol kgw -1 in NaCl, 24 and 2.21 mol kgw -1 in NaClO4. 25 The The stoichiometric coefficients of Mg 2+ are determined by the slope analysis according to the experimental methodology. The 2σ uncertainties resulting from the trapezoidal integration rule were evaluated in the values of log10R and included in the linear fits by weighted least-squares regression analysis. unpredictable results on conditional stability constants, especially for the investigated charge neutralization during which the highly charged ions could interact significantly. 50,51 This problematic situation was avoided by restricting the maximum concentration of Mg 2+ within about 10% of actual ionic strengths at the end of complexation in most series (c.f. Table IV -S2 of the SI). Even though at relatively lower Im, i.e., 0.10 and 0.51 m NaClO4 and 0.10 m NaCl, the last two or three samples exceed somewhat the limit of 10%, the including of these points did not give the second slope (Fig. IV -1). For this reason, they were retained in the slope analysis.

luminescence
Table IV -

Table IV -4 Values of log10[Mg 2+

] and pH obtained from the slope analysis in Fig. IV -1. coefficients are evaluated to be ε(MgUO2(CO3)3 2-,Na + ) = 0.09 ± 0.04 kgw mol -1 in NaClO4 and ε(MgUO2(CO3)3 2-,Na + ) = 0.31 ± 0.02 kgw mol -1 in NaCl. However, from a theoretical point of view it seems difficult to understand that the extrapolation to infinite dilution, i.e. in the absence of salt, would depend on the type of background electrolyte.

Im (NaClO4) MgUO2(CO3)3 2-(log10 [Mg 2+ ], pH) Im (NaCl) MgUO2(CO3)3 2-(log10 [Mg 2+ ],

Discussion on the values of log10β°1.1.3 and ε

The value of log10β°(MgUO2(CO3)3 2-), the Gibbs energy of formation, and specific ion interaction coefficients are reported in the Table IV -1, 2 and 3, respectively. It can be seen that the values proposed in Grenthe, et al. 23 are very similar to the ones obtained here. Our data set presenting the advantage to provide the specific ion interaction coefficients, which could not be selected at the time of the review. The variations of log10β1.1.3 + 16DH vs. Im for MgUO2(CO3)3 2-and CaUO2(CO3)3 2-in NaCl and NaClO4 exhibited virtually identical behaviours and resulted in comparable values of interaction coefficients. In NaCl, ε(CaUO2(CO3)3 2-, Na + ) and ε(MgUO2(CO3)3 2-, Na + ) were determined to be 0.29 ± 0.11 kgw mol -1 and 0.31 ± 0.02 kgw mol -1 , respectively, and in NaClO4, ε(CaUO2(CO3)3 2-, Na + ) and ε(MgUO2(CO3)3 2-, Na + ) were estimated to be 0.09 ± 0.04 kgw mol -1 and 0.02 ± 0.04 kgw mol -1 , respectively. Regardless of the type of alkaline earth cation, the ε values both presented obvious decrease when changing the supporting electrolyte from NaCl to NaClO4. The behaviour of ε implies that the influence of ionic strength on the complexations of Mg 2+ and Ca 2+ with UO2(CO3)3 4-is more sensitively reflected by the change of medium anions than alkaline earth metal ions. At high electrolyte concentrations, the complexation of Mg 2+ and Ca 2+ with UO2(CO3)3 4-can be in all likelihood competed with cationic Na + as well as hampered by attraction of anionic ClO4 -or Cl -through ionic pairings. As a result, the binding processes of Mg 2+ and Ca 2+ with UO2(CO3)3 4-should not be estimated to the binary stage alone. Still, the ion-pairs formations between Mg 2+ /Ca 2+ and the medium anion are supposed to have comparable quantitative effects on the complexation of Mg 2+ /Ca 2+ with UO2(CO3)3 4-, resulting in the close values of interaction coefficients for MUO2(CO3)3 2-in a given electrolyte.

For small ions of high charge density like M 2+ (M = Mg, Ca), the ease of desolvation in their hydrated shells due to the attraction of large ions of small charge density like ClO4 -or Cl -is probably responsible for their ion-pairs association. 26,52 Nevertheless, the ionic attraction could not to overcome the first hydration shell of M 2+ (M = Mg, Ca) in the investigated range of electrolyte concentrations, and thus solvent-shared ion pairs are the most probable type where only one layer of water molecules between the ion partners of the pair exists. 53 Such ion-pairs formation makes the cation M 2+ (M = Mg, Ca) effectively larger, also reduces their electric ability to approach and react with UO2(CO3)3 4-. 54 Whereas once complexed with UO2(CO3)3 4-, the integrity of inner sphere water molecules of alkaline earth cations is disrupted by partly replacing oxygen atoms of carbonates for water molecules. [28][29][30]55 It is worth noting here that complexation of the first and the second alkaline earth cation exhibit slightly different configurations in molecular dynamic (MD) simulations. Under present experimental concentrations, i.e. 0.1-2.21 mol kgw -1 , Ca 2+ -IRVII = 1.06 Å, in Shannon 56 is most probably surrounded with 7 H2O in its immediate environment giving a pentagonal bipyramid geometry 31,33,57 . While smaller Mg 2+ -IRVI = 0.72 Å, Shannon 56 adopts an octahedral structure with perpendicular square plan, which is disadvantageous to the water exchange when coordinating to UO2(CO3)3 4-. 34,58 Kerisit and Liu 33 reported the structural properties of MnUO2(CO3)3 (4-2n)-(M = Mg, Ca, Sr, n = {1;2}) using the MD simulations, and observed that for MgUO2(CO3)3 2-and MnUO2(CO3)3 (4-2n)-(M = Ca, Sr, n = {1;2}) species, at least two water molecules were substituted by carbonates in the first coordination shell of alkaline earth metal ions that promote stable complexation. However, for the Mg2UO2(CO3)3(aq) complex, one of two Mg 2+ exceptionally changed its configuration from bidentate to monodentate towards UO2(CO3)3 4-because of an additional H2O in the inner sphere of Mg 2+ . The intrinsic difficulty in the second Mg 2+ complexation, induced by the solid square geometry in its equatorial plane, possibly contributes to this monodentate binding configuration, also the reason for the non-detection of Mg2UO2(CO3)3(aq) complex in this work. It can also be reminded that Oher, et al. 47 Another important evidence supporting the steric congestion of Mg 2+ has to do with the energy balance obtained for the complexations of Mg 2+ and Ca 2+ with UO2(CO3)3 4-. Indeed, the energy required to break the hydrogen bonds Mg-O(H2O) is higher than the energy released by the formation of Mg-O(CO3) bonds, leading to a net loss of energy. 17 The opposite situation instead applies for the Ca 2+ complexation that gives rise to an exothermic reaction. 16,18 In Fig. IV -S1 of the SI, the clearly discernible linear relationship between the formation constants of Mg and those of Ca with different ligands indicates that the highly charged counterparts generally form stable complexes (see Table IV -S4 of the SI for formation constants). However, with the same ligand, the stability constants pattern is less clear and exhibits two kinds of behaviour depending on the size of alkaline earth metal ions from Ca 2+ to Ba 2+ . For ligands of large charge density and anions of weak acids, e.g. F -, OH -, and CO3 2-, the values of log10Kº decreases generally with the ionic radius. On the contrary, for anions of strong oxoacids, e.g. SO4 2-, S2O3 2-, the stability constants are usually small but present an increasing trend with the cation size (see Table IV -S5 of the SI). Interestingly, the evolution of log10β°1.1.3

shows an unambiguous increase with decreasing size of cation, except the anomaly with Mg 2+

as shown in Fig.

IV -S1 and S2 of the SI. It is therefore suggested that the UO2(CO3)3 4-displays similar characteristics to weak acids in complexation. spectra, irrespective of the nature of background electrolyte or ionic strength. This observation is consistent with the spectral results obtained for Ca-UO2-CO3 complexes in NaCl and NaClO4 in our previous works. 24,25 Nonetheless, the use of 300 lines mm -1 grating is very limited to discriminate the spectral differences of Mg-UO2-CO3 and Ca-UO2-CO3 complexes because of their highly similar spectral fingerprints, and of the convolution of the bands with the optical parameters of the apparatus. In order to probe the possible spectral modifications, we specifically selected two samples at Im = 0.1 and 1 m NaCl to focus on their specific characteristic peaks using the more resolutive 1800 lines mm -1 grating, as done in Oher, et al. 47 The wavelengths corresponding to maximum intensity were found at 465.63 nm, 484.84 nm, 504.66 nm, 527.61 nm and 548.65 nm, as shown in Fig. IV -5. The peak shifts were not such apparent compared with the averaged peak

Evolution of luminescence spectra and temporal properties

The decay-times of MgUO2(CO3)3 2-are estimated by averaging the τ values of correspondent samples that give the slope of one in the slope analysis. Table IV -7 lists the decay-times determined for MnUO2(CO3)3 (4-2n)-(M = Mg, Ca) complexes in NaCl and NaClO4. The possible quenching effect of Cl -on temporal properties of MnUO2(CO3)3 (4-2n)-complexes can be traced by comparing the decay-times measured at 0.1 mol kgw -1 NaCl and NaClO4 (see Table IV -7).

At higher NaCl concentrations, a notable decrease of decay-times is measured for MgUO2(CO3)3 2-as seen for CanUO2(CO3)3 (4-2n)-complexes in Shang and Reiller. 24 shows that the characteristic decay-times decreased steeply in the investigated Im range of NaCl, but varies around an average value of 22.6 ± 1.4 ns in NaClO4, which is in relative good agreement with the values of 18.2 ± 2.7 ns reported by Geipel, et al. 12 and 17 ± 2 ns by Lee, et al. 10 These phenomena are significantly lower than previous results reported for CanUO2(CO3)3 (4-2n)-, with the average decay-times of 34.2 ± 2.2 ns for CaUO2(CO3)3 2-and 59.2 ± 4.2 ns for Ca2UO2(CO3)3(aq) in NaClO4. 25 Table IV -7 Decay-times determined for MnUO2(CO3)3 (4-2n)-(M = Mg, Ca) complexes in NaCl and NaClO4. Ca-UO2-CO3 complexes have longer decay-times than Mg-UO2-CO3 complexes under similar conditions is probably related to the differences in their luminescent capacity.

IV -5 Practical applications

Case of radioactive waste management

In Shang and Reiller, 24 the application of the experimentally measured log10β° and ε values of CanUO2(CO3)3 (4-2n)-was done in the case of water compositions of underground research laboratory for nuclear disposal facilities: the Boom clay rock in Belgium, 36 and the Callovo-Oxfordian clay rock in France. 37 The predictive calculations used the thermodynamic constants of Mg 2+ , Sr 2+ complexes from Dong and Brooks 14 with no ε valuesi.e. ε = 0 using the ThermoChimie 9b database for PhreeqC 43,44 at the time of writing of Shang and Reiller. 24 In this exercise, the ThermoChimie 10a database was used to theoretical speciation calculations. We implement the input PhreeqC file with the log10β°n.1. c,d) clay equilibrium waters using thermodynamic data from the Thermochime 10a database, implementing SrUO2(CO3)3 2-, 14 correcting the MgUO2(CO3)3 2-and CanUO2(CO3)3 (4-2n)-data with thermodynamic constants values from determined in this work and our earlier work, 24 and implementing specific ion interaction coefficient from Table IV -3: (a,b) P(O2) = 10 -68.67 atm and (c,d) 10 -65.83 atmsee Shang and Reiller 24 for details.

Case of seawater

The second practical application to seawater in Shang and Reiller 24 is re-examined here with the newly measured log10β° and ε values for MgUO2(CO3)3 2-. Estimation of uranium speciation in seawater where provided by Maloubier, et al. 59 using Dong and Brooks 14 constants and probably simplified Debye-Hückel ionic strength correctionand Lee, et al., 10 using their determined log10β° extrapolated using the ε(MgUO2(CO3)3 2-,Na + ) = ε(UO2(CO3)2 2-,Na + ) hypothesis. The seawater composition, proposed by Millero, et al. 38 is described in the Table IV -S6 of the SI. The initial seawater information, collected from upper ocean, is known to be supersaturated with respect to calcium and magnesium. 60 The natural concentration of U(VI) seems to be completely soluble under the seawater conditions and not controlled by any common uranyl carbonate minerals of alkaline or alkaline earth metals with the general formula NajCakMgm(UO2)(CO3)3•xH2O(cr)i.e. liebigite, bayleyite, swartzite or andersonite, see Alwan and Williams, 61 Endrizzi, et al., 62 Lee, et al., 63 and references therein. 5 atm from the total carbonate concentration, and using the thermodynamic data from the ThermoChimie 10a database, correcting the MgUO2(CO3)3 2-, CanUO2(CO3)3 (4-2n)-data with thermodynamic constants determined in this work (Mg) and our earlier work. 24,25 We particularly performed calculations with the thermodynamic constants published earlier by Dong and Brooks, 14,15 and Lee and Yun of the SI). This discrepancy can be attributed to the lower value of log10β°(MgUO2(CO3)3 2-) = 25.02 and more negative interaction coefficient (-3 kgw mol -1 ) proposed in Dong and Brooks. 15 The use of these datameasured in NaNO3 within the range of Im = 0.1 mol kgw -1 to 0.5 mol kgw -1needs to be more carefully considered in high NaCl salinity environments.

Building on this exercise, the thermodynamic modelling of seawater contributes to predict and evaluate the competition of magnesium and calcium with triscarbonatouranyl. The distributions of (M-)UO2-CO3 species confirm the predominance of both the -2 bearing charge MUO2(CO3)3 2-regardless of the source of data: log10β°n.1.3 and ε. Even through apparently showing low values, [16][17][18] it has been noticed that enthalpies of the studied complexations are needed in an effort to provide a thoughtful and rigorous evaluation of uranium in the ocean depending on the varying temperature. 64 The results in this work are expected to help addressing the issues in extraction of uranium from seawaters in which Ca2UO2(CO3)3(aq) was conventionally considered to exceed CaUO2(CO3)3 2-. 18,59,62,65,66 It is noteworthy to remember that Maloubier, et al. 59 questioned the result of their speciation calculation with regards to their determination of a too low decay-time results in TRLS in synthetic seawater, which was not in agreement with a predominance of Ca2UO2(CO3)3(aq). From the evolution of decay time of the different complexes as a function of NaCl concentration in Shang and Reiller, 24 as well as in this work, the theoretical speciation and luminescence decay-times are now more consistent.

Case of Canadian aquitards

Another application of the obtained thermodynamic constants is connected with the aqueous speciation calculations in near-surface environments. The vertical variability over the depth of 2.3 to 29.0 meters in pore water compositions representative of smectite-rich clay aquitards in Saskatchewan (Canada) was reported by Hendry and Wassenaar 40 then specifically examined for the eventual uranium(VI)-natural organic matter associations (U(VI)-NOM) by Ranville, et al. 39 The moderately basic and sulphate-type media exhibit regular depth profiles in terms of ionic strength, redox potential, ratio Mg/Ca, U concentration and total carbon (Fig. IV -S4 of the SI).

The oxidized and fractured upper 3 meters of the till is more highly saline with elevated solute concentrations of total dissolved solids (TDS), SO4 2-, Mg 2+ , Na + , and K + because of geochemical weathering. Essentially, the important SO4 2-concentration has been attributed to the oxidation of pyrite occurring at the shallower soil surface (< 3 meters) and the high Na + concentration is sourced from the favoured cation exchange of Ca 2+ and Mg 2+ for Na + after the dissolution of calcite and dolomite in clay minerals. Below 3 meters depth, the till is characterized as an unoxidized and non-fractured zone where the alkalinity decreases with depth until attaining the background concentration below 15 meters depth. The change of oxidation is also reflected by the fact that the total U concentration is reduced by nearly two orders of magnitude from 2.53 µmol L -1 at 2.3 meters depth to 0.084 µmol L -1 at 29.0 meters depth. Particularly, the dissolved salts (TDS, SO4 2-, Mg 2+ , Na + and K + ) display decreasing trends with depth, in contrast to the increasing trend for Ca 2+ . 40 The resulting vertically decreasing ratio Mg/Ca implies an alteration of Mg and Ca speciation with depth so that the reasonable evaluation and the use of appropriate thermodynamic constants are necessary to make justifiable predictions of the real situations.

Reiller, et al. 67 confirmed the negligible proportions of U(VI)-NOM complexes at whatever depth in pore waters by theoretical calculations, but using the Davies 35 Given that the actual temperature of the site is much lower than 25°C, the implementation of ∆fH°m values would be appropriate to estimate the variation of log10β (MnUO2(CO3)3 (4-2n)-) (M = Mg, Ca, n = {1; 2})see Figures S8-1 to S8-8 in the SI in Reiller and Descostes. 7 Because of scarce availability in literature, we decided to include the average of ∆rH°m values determined by Endrizzi and Rao 18 for the Ca-ternary complexeswhich value has been selected by the latter NEA-OECD review 23 and Jo et al. 16,17 for the Mg-and Ca-ternary complexes.

Previously, the particular case of Canadian aquitard has been tested by Reiller and Descostes 7 using PRODATA databasededicated to uranium mining and environmental monitoringwith MnUO2(CO3)3 (4-2n)-(M = Ca, Ba, n = {1;2}, M = Sr, n = 1) from Dong and Brooks 14 and MgUO2(CO3)3 2-from Dong and Brooks 15 and with the ∆rH°m reported by Endrizzi and Rao. 18 We specifically recalled the speciation calculation with PRODATA 1. 7 The values for ∆rHm were listed in Table IV -S8 of the SI.

results were consistent with their prediction of predominant CanUO2(CO3)3 (4-2n)-under these conditions. Despite this, awareness is raised to the fact that the formation constant of UO2(CO3)3 4-in llnl.datapparently chosen from the first NEA-OECD review 21 is not the value selected by Dong and Brooks 14 in their estimation of cumulative constant from successive constants from the first NEA-OECD update. 22 Therefore, the successive formation constants log10K°n.1.3 of MnUO2(CO3)3 (4-2n)-species should have been integrated instead of the cumulative formation constants log10β°n.1.3 using the llnl database. It is also noteworthy that the implemented activity corrections in the parameter file of llnl.datthe extended Debye-Hückel equation with fitted b-dot parameters 76 are not the SIT model used either in NEA-OECD review [21][22][23] or in ThermoChimie 10a database, 45 with which we are conducting the calculations.

For these reasons, we selected specific Mg-containing water information for five wells (see detailed information in Table IV -S9 of the SI) and performed speciation tests: (i) using llnl database with successive log10K°n.1.3 for MnUO2(CO3)3 (4-2n)-from Dong and Brooks; 14 and (ii)

using ThermoChimie 10a database implemented with the cumulative log10β°n.1.3 from the same work. 14 The speciation results displayed only minor even negligible differences (see The water chemistry in North-Western India suggests that aqueous Mg/Ca uranyl carbonate complexes control the U(VI) distribution. However, the research to date on the bioavailability and toxicity of ternary MnUO2(CO3)3 (4-2n)-complexes has tended to only focus on calcium rather than magnesium. For example, Carrière et al. 77,78 proposed negligible biological effects of CanUO2(CO3)3 (4-2n)-complexes on cells. Prat, et al. 79 reported that no clinical manifestations were obtained among the population exposed to drinking water of high concentrations of natural uranium in Finland, when the importance of MnUO2(CO3)3 (4-2n)-was evidenced experimentally and by theoretical speciation. 7,79 In plants, the uptake of uranium was also proven inhibited in calcium-carbonate solutions. 80 The actual uranium toxicity of these samples in India could be studied in light of the MnUO2(CO3)3 (4-2n)-species-controlled distribution of U(VI). estimates with temperature are discussed. The uranium speciation in the case of the safety of nuclear waste management, using the present thermodynamic functions, provides support to the assessment of underground nuclear waste repositories.

IV -8 Supporting Information

V -2 Introduction

The migration of actinides in natural environments, especially uranium, is of great importance in the long-term performance assessment of nuclear waste processing and disposal. 1 It is thus necessary to acquire the environmental information of the geological repository temperature, pH value, nature of solid phases, types of existing inorganic and organic complexing agents, etc. -, which influences the mobility of radionuclides. With this aim in mind, the thermodynamic constantsformation constant log10K°-, and functions -Gibbs energy ΔrG°m, enthalpy ΔrH°m, entropy ΔrS°m, heat capacity ΔrCp°m, etc.for involved reactions are mandatory for the predictive calculations. The complexation of actinides by major inorganic materials in underground waters at other temperatures than standard raises critical scientific interest, because the temperature in the nuclear waste storage casks could be up to 90°C because of radioactive decay. 2,3 However, a certain lack of knowledge on thermodynamic parameters of actinide complexation with temperature renders it difficult to predict chemical behaviour of actinides in the near-field of radioactive waste repositories, where a significant temperature gradient is estimated in timescale evolution. 4 The variation of water properties with temperature, e.g., ion product and dielectric constants, can result in significant rearrangement of the solvation sphere of complexing ions, affecting the energetics of uranium complexation reactions. 5,6 Therefore, full thermodynamic data of uranium reactions with various common ligands like carbonate are of interest to accurately simulate the uranium transport in the nearfield, as well as in the far field, of a repository. 7 The nuclear waste repository isolation system consists typically of a natural barrier, mainly in clay or granite, depending on the country-specific rock formation, and an engineered barrier including in general waste canister, bentonite backfill, and concrete liner. 8 This multi-barrier system is extremely important in preventing the migration of actinides into the biosphere by means of pore-waters, which are similar in compositions with Cl -, SO4 2-, HCO3 -, Ca 2+ , Mg 2+ , and Na + as the main ions at pH value ca. 8 in several envisaged geological formations. [9][10][11][12][13][14] The triscarbonatouranyl(VI) complexes with alkaline earth metals (Mg 2+ , Ca 2+ ) have been reported, or calculated, in the literature as one of the predominant U(VI) species under a variety of conditions at near-standard temperatures, [15][16][17][18][19][20] even if U(IV) phases are controlling the to extrapolate the log10K°n.1.3 to 70°C and derive the enthalpies using the Van't Hoff equation.

The enthalpies of reaction were determined to be ∆rH°m = -1.2 kJ mol -1 and -1.7 kJ mol -1 for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), respectively, in satisfactory agreement with the results directly measured at 25°C by calorimetry. The application of isoelectric reaction to enlarge the validity of the Van't Hoff equation is particularly useful to predict the stability constants at nonstandard temperature. The authors later applied the same analysis method to Mg-UO2-CO3 system and obtained ∆rH°m = (4.3 ± 1.5) kJ mol -1 for MgUO2(CO3)3 2-. 29 Recently, Maia, et al. 40 also using the Van't Hoff equation and a Schubert type method applied to an ion-exchange resin, reported ∆rH°m = (14.9 ± 5.1) kJ mol -1 and (36.5 ± 10.0) kJ mol -1 for CaUO2(CO3)3 2-for Ca2UO2(CO3)3(aq), respectively, slightly more important than the previous results. 18,28 One should question before any modelling the fulfilment of an important criterion of the Van't Hoff equation, the constant molar enthalpy in the investigated temperature range, as the derived parameters are very sensitive to the fitting values. Maia, et al. 40 also reviewed the data of Jo, et al. 28 by directly fitting the experimental values measured at 10-70°C to the Van't Hoff equationnot adopted in the original paper. Similar thermodynamic trends were obtained with the data of Jo, et al. 28 -∆rH°m = (10.9 ± 9.5) kJ mol -1 and (23.5 ± 7.0) kJ mol -1 for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), respectivelythat could be attributed to the constant enthalpy assumption in absence of modification to isoelectric reactions.

In the present work, we will focus on the complexation of triscarbonatouranyl(VI) with Ca 2+

and Mg 2+ by TRLFS in the temperature range 10-50°C and 5-30°C, respectively. Stability constants of MnUO2(CO3)3 (4-2n)-will be determined at Im = 0.1 mol kgw -1 NaCl, with spectroluminescent titration by Ca 2+ or Mg 2+ at variable pH values, which has been successfully applied to determine formation constants and interaction coefficients in our previous studies. 34- 36 The luminescence emission intensity and decay-times with increasing temperatures will also be studied. Different approximation models -the Van't Hoff equation, the constant heat capacity model, the DQUANT equation, and the isoelectric modificationwill be adjusted to our experimental values. The molar enthalpy values obtained in different approaches will be compared with previously published values. 18,28,29,40 The implications in complexation mechanisms will be further discussed, in combination with our previous observations in concentrated electrolytes. These data can serve as parameters in the prediction of uranium migration in deep underground geological nuclear waste management context. Coupled at the desired temperature in order to maintain the solution temperature during pH measurements. The pH values were regularly adjusted if there were any modifications prior to the luminescent spectrum acquisition. The pH measurements were performed in triplicate for each sample.

Time-resolved laser-induced luminescence spectroscopy

The time-resolved luminescence apparatus has been already described. [34][35][36] The sample solution was added in a fluorescence quartz cuvette with a sealable cap (QS 117F, Suprasil, Hellma Analytics) leaving a minimum headspace to limit evaporation and CO2 degassing.

D τ i ) i (1 -exp (- W τ i )) (V -1)
where FI and FI0 are respectively the fluorescence intensity at a specific delay time and at D = 0.

The gate width W = 1 μs and τi represents the decay-time of species i.

interaction coefficients from Table V -S4 of the SI. 34,36,38 As the H + activity was determined from the acidity measurement of the solutions, the dissociation constants of water pKw, were used to evaluate the H + concentration at various temperatures. 5 The corresponding α values at each specified aqueous condition was listed in Table V -S5 of the SI.

The specific ion interaction theory (SIT) was used to correct the ionic strength effect in solution, as recommended in widely accepted Nuclear Energy Agency-Thermochemical Data Base (NEA-TDB) reviews. [37][38][39] The activity coefficient γi of an aqueous species i in the SIT approach is given by

log 10 γ i = -z j 2 D H + ε(j, k, I m ) m k (V -5)
with DH(T) the temperature-dependent Debye-Hückel term expressed as,

D H (T) = A(T)√I m 1+1.5√I m (V -6)
where zi is the charge of ion i, ε(j,k,Im) (kgw mol -1 ) is the ion interaction coefficient between the aqueous ion j and ion k of background electrolyte of concentration Im at a given temperature in this study no variation of ε with Im was postulated and is thus considered ε(j,k) in the following. The temperature influence on ε(j,k) (kgw mol -1 ) is usually not available for the involved ions. However theoretically, ε(j,k)(T) = ε(j,k)(T°)T°/T because ε(j,k)mk is a second virial expansion term. 45,46 In the experimental temperature range, the correction of ε(j,k)mk to temperature is not necessary as the correction should be too small to be considered. where

Δz 2 = z(M n UO 2 (CO 3 ) 3 (4-2n)-) 2 -n z(M 2+ ) 2 -z(UO 2 (CO 3 ) 3 4-) 2 (V -8)
(∆z 2 = -16 for n = 1 and ∆z 2 = -24 for n = 2) and

Δε = ε(M n UO 2 (CO 3 ) 3 (4-2n)-, Na + ) -nε(M 2+ , Cl -) -ε(UO 2 (CO 3 ) 3 4-,Cl -) (V -9)
with the used values of ε(j,k) listed in Table V -S4 of the SI.

Temperature dependence modelling

The formation constants for many chemical species are tightly related to temperature.

However, thorough understanding of the temperature-dependent chemical behaviour of MnUO2(CO3)3 (4-2n)-complexes has not yet been developed because of the scarcity of experimental results at higher than standard temperatures. 39 Nonetheless, computational methods such as the constant enthalpy equation, the constant heat capacity equation, the DQUANT model and the modification to isoelectric reactions have been proven adequate in a certain temperature range for evaluating temperature-dependent complexation constants. These Rln( 10)

( 1 T - 1 T 0 ) (V -10)
with K° the formation constant under standard conditions (infinite dilution), R the gas constant (8.31451 J K -1 mol -1 ), T the absolute temperature in Kelvin, and ∆rH°m(T0) the molar reaction enthalpy at standard temperature T0. It should be noted that the validity of temperature range (T-T0) ≤ 10 K is quite narrow in this equation. 47 Extrapolation over a temperature range larger than 10 K will undoubtedly introduce greater uncertainties in the resulted log10K°(T).

The constant non-nil heat capacity equation slightly extends the validity of temperature range to 20 K (Eq. V -11). 47 The assumption of constant heat capacity is useful to calculate the equilibrium constant in the range 273-473 K. In some cases, the heat capacity term can be neglected if the entropy and the heat capacity are comparable in magnitude as the term of ∆rC°p,m[∆T-T ln(T/T0)] is less important than ∆rS°m∆T. log 10 K°(T)= log 10 K°(T 0 ) -

∆ r H m °(T 0 )
Rln( 10)

( 1 T - 1 T 0 ) + ∆ r C p,m °Rln (10)
[(

T 0 T -1) +ln(T/T 0 )] (V -11)
Helgeson 48 reported the enthalpy-and entropy-based DQUANT model, as shown in Eq. V -12.

log 10 K°(T)= ∆ r S m ∘ RT ln(10) {T 0 - θ ω [1-exp (exp(b+aT) -C+ T-T 0 θ )]} - Δ r H m ∘ (T 0 )
RT ln (10) (V -12)

where ∆rH°m represents the molal reaction enthalpy, ∆rS°m denotes the molal reaction entropy, a = 0.01875 K -1 , b = 12.741, c = exp (b + aT0), θ =219 K and ω = (1 + acθ). The uncertainties of the formation constants determined for various chemical reactions using the DQUANT equation show less than 1-9% error in 0-200°C temperature range. 47 Isoelectric reactions are defined for the chemical reactions of the following form: the sum of positive charges among the reactants equals the total amount of positive charges among the products, and the same for the negative charges among reactants and products. The modifications of MnUO2(CO3)3 (4-2n)-complexation to isoelectric reactions requires the additional equations Eq. V -13 and Eq. V -14 to balance out the charges in Eq. V -2:

CO 2 (aq) + H 2 O ⇄ CO 3 2-+ 2H + (V -13) UO 2 2+ + 3CO 3 2-⇄ UO 2 (CO 3 ) 3 4- (V -14)
The isoelectric forms of MnUO2(CO3)3 (4-2n)-complexation are expressed as follows.

M 2+ + UO 2 2+ + 2CO 2 (aq) + CO 3 2-+ 2H 2 O ⇄ MUO 2 (CO 3 ) 3 2-+ 4H + (V -15) 2M 2+ + UO 2 2+ + 3CO 2 (aq) + 3H 2 O ⇄ M 2 UO 2 (CO 3 ) 3 (aq) + 6H + (V -16)
These kinds of reactions can cancel out the electrostatic contributions to the temperature dependence of enthalpy. The formation constants of the additional reactions Eq. V -13 and V -14 have been calculated for the temperature range of 5-90°C using the existing general formula, [49][50][51] as listed in Table V -S3 of the SI. The enthalpies for the isoelectric reactions Eq.

V -15 and V -16 were calculated by the Van't Hoff equation with the assumption of constant enthalpy and zero heat capacity. In this study, the formation constants log10K°n.1.3(T) of MnUO2(CO3)3 (4-2n)-were finally determined by subtracting the log10K°(T) of Eq. V -13 and Eq.

V -14 from the global equilibrium constants log10β°n.1.3(T) for the isoelectric reactions Eq. V - displays an obvious decrease in intensity when the temperature increases from 40 to 50°C.

V -4 Results and discussion

Luminescent emission spectra analysis

Apparently, the increase in temperature leads to slight shifts of the spectra to higher wavelengths also broader characteristic bands. The band shifts and broadening can become more significant at high temperatures even contributing to the deformation of spectra. One can observe that in the temperature range of 10-35°C, the samples exhibited well-defined characteristic spectra in the course of Ca-UO2-CO3 complexation, while outside this range, the spectra showed irregular evolution at the beginning of the complexation, i.e., 40, 45, and 50°C.

Specifically, we compared the area-normalized emission spectra of triscarbonatouranyl from 5 to 50°C in the absence of Mg 2+ /Ca 2+ in Fig. V -4. The concentration of U(VI) varied from 50 to 250 μmol kgw -1 in order to increase the luminescence at the highest temperatures. Indeed, the emission spectra contain features originating from vibronic transitions from the lowest excited state to the ground state manifold, 11→00 of the hot band and 10→0ν (ν = 0, 1, 2, 3) of other peaks. 52 With the increase in temperature, the less-defined bands of UO2(CO3)3 4-were continuously red-shifted. impression that the non-radiative decay becomes more important at higher temperature, which lowers the quantum yield (Q). It is also possible that a change in temperature will alter the rates of some reactions, e.g., the formation of an internal charge-transfer (ICT) states, which hinders the radiative decay. 53 Further research is needed to probe the luminescent properties of triscarbonatouranyl(VI) with temperature.

The evolution of spectral band positions with increasing [Ca 2+ ] allows sorting out individual spectra of involved species, notably at higher temperatures, i.e., 40, 45, and 50°C, where the band shifts were obvious. The data analysis method employed in the present study is the combination of Singular Value Decomposition (SVD), 54 used to assess the number of distinguishable components, and multivariate curve resolution alternating least squares (MCR-ALS) [55][56][57]

F(λ, [Ca])= ∑ F k (λ)C k ([Ca]) k (V -17)
where F(λ, [Ca]) (luminescent intensity), represents an additive result of pure spectra Fk(λ), of k principal components determined by SVD weighted by their concentration profile as a function of Ck ([Ca]). This interpretation of luminescent spectra evolution is also chemically meaningful. However, it was noted in Ruckebusch, et al. 55 that SVD can suffer from an excess of estimated rank towards the expected number of components when the pure spectrum of a single chemical compound shows a clear spectral shift, thus the sequential transformation of discrete components was applied in the decomposition in order to verify the concentrationwavelength separability. 58 The global matrices data analysis was conducted for the Ca series at 50°C. The first spectrum [Ca 2+ ] = 0 was separately implemented as an external file serving as the reference shape, which can facilitate the calculation and improve the accuracy. The SVD analysis and the sequential transformation were tested for determining the number of components. The spectra and concentration profile were constrained to be positive. It was shown by both methods that there are two distinguishable components, which is consistent with the hypothesis of CaUO2(CO3)3 One can observe that at 50°C, the energy losses in luminescence are intensified for the incomplete complexations, i.e., UO2(CO3)3 4-and CaUO2(CO3)3 2-, that contribute to the redshifted and strongly deformed spectra. On the contrary, the neutral charged species Ca2UO2(CO3)3(aq) is more resistant to the temperature effect. The characteristic band positions were determined at 465, 484, 505, 526, and 549 nm for Ca2UO2(CO3)3(aq) that are consistent with our previous results, 34,35 whereas the pure spectrum of CaUO2(CO3)3 2-at 50°C exhibited obvious red-shifts with bands resolved at 480, 497, 515, 533, and 552 nm. The recovered spectra using the MCR-ALS analysis of CanUO2(CO3)3 (4-2n)-at 40 and 45°C are not shown, but the trends in the spectral features are similar. The same procedure was also applied to the Mg-series at 30°C. As shown in Fig. V -S3 of the SI, the pure spectrum of MgUO2(CO3)3 2-at 30°C did not show a shift compared to our previous measurements, 36 with band positions determined at 466, 485, 503, 523, and 546 nm. It is believed that the spectral properties of Mg/Ca-UO2-CO3 complexes reflect subtle differences in the coordination of alkaline metal ion to triscarbonatouranyl(VI), which also affect their thermodynamic properties. This will be pointed out with the analysis of their thermodynamic functions in the following section.

Determination of enthalpy and entropy

The enthalpy of the formation of MnUO2(CO3)3 (4-2n)-is the net thermal effect of three processes: i) the energy required for dehydration of Mg 2+ /Ca 2+ and UO2(CO3)3 4-; ii) the energy released/required during the formation of MnUO2(CO3)3 (4-2n)-; and iii) the energy released/required during the formation of hydrogen bonding between dehydrated water molecules and the surrounding bulk water. The value of ∆rH°m is established by the competition between the energy-consuming and -generating processes in complexation and a net gain/loss in energy balance resulting in negative/positive value of ∆rH°m, respectively. The stability constants of MnUO2(CO3)3 (4-2n)-determined in this work did not show obvious variation in the investigated temperature range that made the model parameters sensible to adjustment and resulted in large errors. Four approximation approaches were tested to fit the experimental log10K°n.1.3: the Van't Hoff equation, the constant heat capacity equation, the DQUANT equation, and the isoelectric modification. Notwithstanding the different assumptions of each estimated model, as well as their corresponding temperature range of validity, the four approximation approaches provide satisfactory fitting of experimental data and allow prediction of the formation constants with temperatures. It must be reminded here that two digits will be kept for thermodynamic constants and entropy values, and three digits will be kept in the values of the Gibbs energy and enthalpy values, in coherence with the selected data of the NEA-TDB, [37][38][39] even if they may not be significant.

With the values of log10K°n.1.3 and ∆rH°m determined at 25°C using different approximation approaches, the temperature effect can be quantitatively estimated to predict the formation constants as a function of temperature. The decrease in entropy term for the first and second stepwise reactions may results principally from the degree of solvation and the polarity of surrounding water molecules of the reacting species and the complexes. The hydration shell of UO2(CO3)3 4-highly charged and symmetric in structureis much more ordered than that of the globally -2 bearing charge ion CaUO2(CO3)3 2-. In the process of complexation, the ordered water molecules around UO2(CO3)3 4-have to be released to a disordered bulk solvent that allows a bidentate binding configuration for calcium. Therefore, the initial hydration shell of UO2(CO3)3 4-is inevitably perturbed by the first complexation, which explained the great entropy gain as a result of a large change in hydration sphere. The molecular calculations showed that the triscarbonatouranyl(VI) group can impose a specific geometry on the primary hydration sphere of calcium ion by reducing its coordination number from 7 to 5. 59 In this respect, the second complexation of Ca 2+

to CaUO2(CO3)3 2-should be less favourable in terms of entropy because the distorted hydration shell of CaUO2(CO3)3 2-may be easier to accept the next Ca 2+ than the first one, which would result in a lower entropy gain. 60 In addition, it was postulated in classical molecular dynamics simulations 61 that the presence of solvent ion Na + is able to distort one of the three equatorial carbonate groups in CaUO2(CO3)3 2-from bidentate to monodentate, while the second Ca 2+ is able to stabilize the bidentate binding of the three CO3 C, IRVI = 0.72 Å in Shannon 62 assuming Mg 2+ six-coordinated by H2O) exhibits a stronger polarizing effect than Ca 2+ (charge denticity of 64 C mm -3 with IRVII = 1.06 Å in Shannon 62 assuming Ca 2+ seven-coordinated by H2O) in aqueous solution. When complexing to large negative ion like UO2(CO3)3 4-the smaller Mg 2+ should distort the electron cloud of UO2(CO3)3 4- to a larger degree, and even could favour partial covalence between Mg and OCO 3 , which leads to a more consequent deformation of originally symmetric structure of UO2(CO3)3 4-. Therefore, MgUO2(CO3)3 2-should be less symmetric in electron distribution and exhibits stronger polarizing power than CaUO2(CO3)3 2-resulting in more ordered surrounding hydration sphere. This is in general agreement with the Hofmeister series, and kosmotrope character of Mg 2+ relative to Ca 2+ . 63 This subtle difference was detected in TRLFS using a more dispersing 1800 lines mm -1 grating. 36,52 In nuclear waste management, the high-level waste (HLW) package disposal facilities are composed of multiple barriers to prevent the leaching of radionuclides from the stainless steel canisters (the primary barrier) and their contact with interstitial water. Theoretically, the thick steel overpacks wrapping the canister can protect the vitrified HLW from water corrosion for 500 years. Once the HLW packages in underground repositories, the temperature of the cell can attain around 80°C in 10-15 years due to the generated heat by radioactive decay, but the temperature between two HLW cells needs a longer period to reach the peak, e.g., 40-60°C after approximately 400-500 years dependent on the types of geological formation and waste. 4 To predict the uranium behaviour under repository conditions at temperatures higher than 25°C, the thermodynamic calculations were performed to determine the uranium speciation and to identify the mineral phases likely to precipitate using Thermochimie 10a database file, 69 in which the uranium chemistry is largely based on NEA-OECD analyses. [37][38][39] The thermodynamic parameters for the interactions of UO2-CO3 with Mg/Ca were based on the isoelectric reaction approximation because of its higher validity range of temperature that would render more reliable estimates. The thermodynamic parameters in the form of analytical expression to implement in the input PhreeqC file are listed in Table V -3. We noted that the determined values of log10K° are so sensitive to the parameters in analytical expression that they should not be rounded off in calculation. Water compositions in the Boom and COx geological formations were taken from de Craen, et al. 9 and Gaucher et al. 11 The aqueous uranium concentration was set in equilibrium with the dissolution of UO2:2H2O(am) phase in both calculationsadapted from UO2(am,hyd) in NEA-OECD analyses 38,39 that the ∆rH°m value is absent from Thermochimie 10a database file. 69 83 atm using thermodynamic data from the Thermochimie 10a database, 69 implementing SrUO2(CO3)3 2-, 26 specific ion interaction coefficients of Mgn/CanUO2(CO3)3 (4-2n)-, [34][35][36] and their ΔrH°m values determined in this work based on the isoelectric reaction approximation.

Table V -3 Parameters in the expression of log10K°n.1.3(T) = A1+A2 T+A3 T -1 +A4 log10T+A5T -2 (T in Kelvin) based on the isoelectric reaction modification in

Hence, the formation of U(OH)4(aq) may thus play an important role in reduced uranium transport. At 40-60°C, our simulation results suggest that uranium would primarily exist in the reduced state in the form of U(OH)4(aq) in the COx and Mol waters and the stability of UO2 can be sustained by the presence of pyrite (FeS2), which consumes the oxygen supplied by infiltrating water. 70 In the very long term, the temperature in the near field of the repository will decrease to ambient temperature. 4 As hard-soft-acid-base principles suggested, the uranyl ion, a hard (high charge to radius ratio) cation, should form strong complexes with carbonate ions, which are hard anions. Thus the uranium migration would be promoted because of CanUO2(CO3)3 (4-2n)-predominance in the COx and Mol waters as suggested by previous calculations. [34][35][36] Furthermore, we note that the variation of uranium solubility as a function of temperature depends on the chosen solubility controlling phase in calculation. There is an underlying assumption in HLW management that reducing conditions will prevent uranium migration due to the formation of uranium ore deposits. [71][72][73][74][75][76][77] The principal ore minerals includes two main typesuraninite and pitchblende, both nominally UO2that differ from each other in their depositional setting and chemical constitution. [78][79][80] As no value of enthalpy for UO2:2H2O(am) selected into the Thermochimie 10a database file, special focus was put on the Mol and COx water systems in equilibrium with the dissolution of uraninite UO2(cr) for which the ∆rH°m value is available.

V -6 Conclusions

The present study reports the formation constants for the complexation of atm using thermodynamic data from the Thermochimie 10a database, 12 implementing SrUO2(CO3)3 2-, 13 specific ion interaction coefficients of Mgn/CanUO2(CO3)3 (4-2n)-, [4][5][6] and their ΔrH°m values determined in this work based on the isoelectric reaction approximation. Al-Batin Alluvial Fan (southern Iraq), 12 and in the Lodève basin (France) 13 ;

(c) waters crossing the Cigar Lake mining site (Saskatchewan, Canada), 14 and in the Oklo uraninite deposit area 15,16 ;

(d) well waters as local sources of drinking sources in Finland, 17 around the Semipalatinsk Nuclear Test Site in Kazakhstan 18 , as well as near-surface Canadian aquitards considered as underlying groundwater 19,20 ;

(e) waters sampled upstream, downstream and within uranium mining sites, e.g., Cluff Lake in Canada 21 , Lac Saint-Clément in France 22 ;

(f) standard seawater 23,24 ;

(g) several waters compositions issued from the general dataset generated through the European project FOREGS (Forum of European Geological Surveys * ) 25 ;

VI -2 Description of studied water compositions

The chemical composition of natural water is a result of the physical, chemical, and biochemical processes that occur in the atmosphere-soil-rock water system. A wide range of chemical elements can dissolve in natural waters due to interactions with the atmosphere, the surrounding environment, soil, and bedrock. In particular, the rock-mineral solubility has a decisive influence on the chemical composition of natural water. Type and amount of major, secondary, and minor elements, as listed in The water compositions are divided into groups according to their value of log10(a(Na + )/a(H + )) and separately presented in the diagram having the nearest value of log10(a(Na + )/a(H + )), as shown in Fig. VI -6. Ca2UO2(CO3)3(aq) can be identified in a large field at high log10(a(HCO3 -) a(H + )) and high log10(a(Ca 2+ )/a(H + ) 2 ) in the predominance diagram with log10(a(Na + )/a(H + )) varying from 2 to 7. Actually, the variable log10(a(Na + )/a(H + )) partially reflects the ionic strength of the water samples as their pH values are comparable and sodium represents one of the major cations in these water samples. That is to say, there is close relationship between ionic strength and log10(a(Na + )/a(H + )) for the selected water solutions. It is worthy to point out that the dominant species calculated for several water samples do not strictly correspond to their position in the diagram because the predominance diagrams are calculated taking only NaCl as background electrolyte.

VI -5 Analysis of results

One can observe that the predominance regions of UO2(CO3)3 4-and CaUO2(CO3)3 2-enlarge with increasing log10(a(Na + )/a(H + )) as a result of increasing interionic attraction between Ca 2+ and other major anions. When viewed as a function of log10(a(Na + )/a(H + )) in the predominance diagram, the field of Ca2UO2(CO3)3(aq) confined by the precipitation of calcite narrows with increasing log10(a(Na + )/a(H + )). TRLFS can be a useful technique to elucidate the speciation of uranium in water because of the very low detection limits, a high sensitivity with respect to U(VI) complex, and its triple selectivityi.e., selectivity in excitation wavelength, emission spectrum, and time-resolved acquisition. In general, TRLFS provides two important spectroscopic information: characteristic luminescence decay time and positions of the emission maxima. In a fingerprinting procedure, the recorded TRLFS spectra can be identified with the help of previously acquired reference complexes. In the preceding chapters, we have obtained these primary spectroscopic information for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes under varying conditionstype of electrolyte, ionic strength and temperaturethat will be compared with the recorded spectra of Lodève water samples.

As previously described, the acquisition were completed with a 355 nm tripled Nd:YAG laser (Surelite, Continiuum, USA) delivering about 170 mJ of energy, pulsed at 10 Hz with 5 ns

pulse. An optical parametric oscillator system (Horizon, Continiuum, USA) is set at λex = 450 nm. At this excitation wavelength, the luminescence quantum yield for uranium(VI) is favourable. The fluorescence signal is focused into a monochromator spectrometer (Acton) using a 300 lines mm -1 grating. Emission spectra are recorded as a function of gate delay D ≥ 25 ns, minimum for preventing the Raman scatter peak of H2O, at the fixed gate width W = 1 μs.

The two natural occurring water samples in Lodève both delivered a measurable luminescence signal, which is characteristic for uranyl carbonate luminescence. present in natural waters can impair the detection of total uranium because of the complexation of organic species, such as humic substances, with uranyl ion. Furthermore, one can expect luminescence emission of these ligands if they contain delocalized p-electron system, like in aromatic compounds. 41 The previous determination of decay times using bi-exponential fits (Fig. VI -12) implies the possible existence of organic ligands in the aqueous system due to the very short decay time of ꚍ1 = (11 ± 1) ns. VII -2 Study of Ca n UO 2 (CO 3 ) 3 (4-2n)-complexes in NaClO 4

Similar study at higher ionic strength was thus performed in NaClO4 medium owing to its less important quenching effect than NaCl on triscarbonatouranyl units. The range of ionic strength was extended to 2.46 mol kgw -1 , while a high [ClO4 -] ca. 3.50 mol kgw -1 could still induce quenching effects towards CanUO2(CO3)3 (4-2n)-complexes. The specific ion interaction coefficients were estimated to be ε(CaUO2(CO3)3 2-,Na + ) = (0.02 ± 0.04) kgw mol -1 and ε(Ca2UO2(CO3)3(aq),NaClO4) = (0.18 ± 0.07) kgw mol -1 . These latter values were different from the ones obtained in NaCl.

Several modified models based on the SIT approach were applied to better represent our experimental data and examine the dependence of interaction coefficients on ionic strength. We estimated a reliable prediction of thermodynamic constants up to 3 mol kgw -1 . In a larger range of ionic strength, the ion associations, e.g., Na + -UO2(CO3)3 4-, Na + -CaUO2(CO3)3 2-, and Ca 2+ -ClO4 -/Cl -, could contribute to increasing ion-pairing formations where the activity coefficient should be estimated by the Pitzer model with better accuracy. This work provides valuable information to address the interaction effects between Ca-UO2-CO3 species and 1:1 type electrolytes. The differences in ∆ε in NaCl and NaClO4 may partly be explained by the ion triplet interactions of Na + -Ca 2+ -ClO4 -/Cl -that could reduce the activity of Ca 2+ to different degrees depending on the kind of anion with increasing ionic strengths. It is suggested that the affinity of the cation in background electrolyte with CanUO2(CO3)3 (4-2n)-(n = {1;2}) are to be taken into consideration at high ionic strengths, especially for the globally non-charged species.

VII -3 Study of MgUO 2 (CO 3 ) 3 2-complex in NaClO 4 and NaCl and applications to different geochemical contexts.

As the log10βn.1.3 vs. Im functions displayed remarkably different patterns for CanUO2(CO3)3 (4-2n)-complexes in NaCl and NaClO4, we proposed new determinations of thermodynamic constants for MgnUO2(CO3)3 (4-2n)-complexes in these two media by the same method. A slope ca. 1 was obtained in the linear correlation at ionic strength 0.1 to 1 mol kgw -1 NaCl and 0.1 to 2.21 mol kgw -1 NaClO4, indicating the formation of MgUO2(CO3)3 2-, while the expected presence of Mg2UO2(CO3)3(aq) cannot be identified.

  Formation des complexes ternaires alcalino-terreux/uranium/carbonate. Données thermodynamiques et influence sur la spéciation et la solubilité de l'uranium dans différents contextes géochimiques Mots clés : Uranium, Métaux alcalino-terreux, Carbonate

Fig. 1

 1 Fig. 1 Diagrams of the distribution of aqueous uranyl species in our solutions experimental, calculated from recommended thermodynamic data

Fig. I - 1

 1 Fig. I -1 Equatorial coordination of U(VI) in the [UO2(CO3)3 4-] complex.

I - 4

 4 Fig. I -2 Jablonski diagram including vibrational levels for absorbance, fluorescence, phosphorescence and non-radiative decay-internal conversion, vibrational relaxation, intersystem crossing.

4. 1 . 2

 12 De-excitation mechanismsImmediately after the absorption of a photon, de-excitation processes will occur with different probabilities. A part of the absorbed energy is dissipated via a non-radioactive pathway to reach the lowest vibrational energy level of the first excited state (S1 = 0; Fig.I -2), during which the excess energy is transferred to the rotational and vibrational degrees of freedom. This process is known as internal conversion or vibrational relaxation (loss of energy in the absence of light emission) and is usually completed in a picosecond or less. This part of the dissipated energy can then be converted into heat, which is absorbed by the surrounding solvent molecules when colliding with the fluorophore in the excited state (dynamic inhibition). Although radiative de-excitation from S2 to S0 is also probable, the corresponding emission is of low energy with a very short decay-time, making detection often inaccessible. The loss of energy through a non-radiative transition is more evident at high temperatures, resulting in larger lines widths in the emission spectra.The excited molecule exists in the lowest excited singlet state (S1 = 0; Fig.I -2) for periods on the order of nanoseconds before relaxing directly to the ground state (fluorescence) or changing to another excited state of spin multiplicity 3 (triplet state T1) by spin change and inter-system crossing (phosphorescence), which is generally unfavorable and occur with longer decay-time. The phosphorescence can result either in photon emission from the triplet state T1 or in transition back to the excited singlet state S1 generating the delayed fluorescence.

  given fluorophore is theoretically affected by multiple interactions with the solvent molecules in their local environment. These interactions can have dramatic effects on emission spectra, such as spectral shifts to longer wavelengths or spectral deformation due to energy losses. The related factors possibly having major effects on the energy of the excited state are illustrated in Fig. I -3.

Fig. I - 3

 3 Fig. I -3 Effects of environment on the energy of the excited state. The dashed arrows indicate the fluorophore can be fluorescent or nonfluorescent in the different states.

Fig. I - 4

 4 Fig. I -4 Jablonski diagram for fluorescence with solvent relaxation

log 10 𝛽

 10 𝑗 °= ∆ log 10 γ + log 10 β 𝑗 (I-21) with ∆ log 10 𝛾 = log 10 𝛾 (𝑀𝐿 𝑗 ) -log 10 𝛾 (𝑀) -𝑗 log 10 𝛾 (𝐿)

2 (I- 33 )

 233 Using the temperature derivatives of enthalpy and entropy, one can obtain:

Fig. I - 8

 8 Fig. I -8 gives the distribution of CO2(aq), HCO3 -and CO3 2-as a function of pH in 0.1 MNaNO3 at P(CO2)= 10 -3.5 atm. Under these conditions, HCO3 -predominates over the range of pH 4-9 with the concentration increasing from about 10 -6 M at pH 4 to 10 -2 M at pH 8 and CO32- is always in the minority compared to HCO3 -at the same pH.

Fig. I - 8

 8 Fig. I -8 Distribution of carbonate species in equilibrium with air (P(CO2)= 10 -3.5 atm) as a function of pH. The thick line represents the total dissolved carbonate concentration. (Electrolyte solution: 0.1 mol kgw -1 NaCl)

  Fig. I -9 shows the speciation diagram of U(VI) at the concentration of 50 µmol kgw -1 in a solution of Im = 0.1 mol kgw -1 NaCl at T = 298 K and in equilibrium with the partial pressure of carbon dioxide of 10 -3.5

Fig. I - 9

 9 Fig. I -9 Speciation diagram of [U(VI)] = 50 µ mol kgw -1 in a solution of Im = 0.1 mol kgw -1 NaCl at T = 298 K, P(CO2) = 10 -3.5 atm.

Fig. I - 10

 10 Fig. I -10 (a) Fluorescence decay spectra of the sample at [U(VI)] = 50 μmol kgw -1 , pH = 7.6 equilibrated with atmospheric CO2(g), recorded with λex = 450 nm from D = 20 ns to D = 210 ns. (b) Red line: spectrum recorded at D = 20 ns. Blue line: spectrum at D = 210 ns. Green line: subtraction of the blue line from the red one, also characteristic spectrum of UO2(CO3)3 4- species.

  )34-+ 13H + ⇄ (UO2)3(OH)5 + + 9CO2(g) + 4H2O log10 * β ⸰'' 3.5.0 = log10 * β ⸰ 3.5.0 -log10β ⸰ 1.0.3 -9log10β ⸰ 0.1.1 log10 * β ⸰'' 3.5.0 = 82.28 ± 0.84

Fig

  Fig. I -11 (a) Luminescence spectra of UO2 2+ in 5% H3PO4 medium at different uranium concentrations from 10 -8 to 10 -10 M, recorded at excitation wavelength of 429 nm. (b) Linear dependence of luminescence peak intensities at 485.4 nm on uranium concentration.

  Fig. I -12 Luminescence spectra collected for uranium concentrations down to 10 -11 M in phosphoric medium.

  The luminescence spectra were collected for uranium concentrations down to 10 -11 mol kgw -1 , as shown in Fig. I -12. It is clear that the characteristic peaks cannot be distinguished at [U(VI)] = 10 -11 mol kgw -1

  𝑔 = 𝐸 𝐴𝑔/𝐴𝑔𝐶𝑙 °+ 𝐴 0 log 10 𝑎 𝐶𝑙 -,𝑔 (I-75) 𝐸 𝑀 = 𝑟𝐴 0 log 10 (𝑎 𝐻 + ,𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 /𝑎 𝐻 + ,𝑔 ) (I-76) 𝐸 𝑗 ∝ 𝑓(𝐼 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ) -𝑓(𝐼 𝑟𝑒𝑓 ) (I-77) where A0 = ln(10)RT/F with F the number of Faraday (A0 = 59.16 mV/pH at 25°C), and r is a coefficient generally close to 1 that measures the deviation of the membrane response form the theoretical response. The activities aCl-,ref , aCl-,g , and aH+,ref are intrinsic parameters that are related to the electrode solutions and normally known to users. The liquid junction potential Ej can be approximately described by a certain function f, which will be evaluated further for different reference filling solutions. With these variables and approximations, the measured potential can be written in the following form. 𝐸 𝑚𝑒𝑠 = 𝑟𝐴 0 log 10 𝑎 𝐻 + ,𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 + 𝑓(𝐼 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ) + 𝐵 + 𝐶 (I-78) with 𝐵 = 𝐴 0 log 10 𝑎 𝐶𝑙 -,𝑟𝑒𝑓 -𝑓(𝐼 𝑟𝑒𝑓 ) (I-79) 𝐶 = -𝐴 0 (𝑟 log 10 𝑎 𝐻 + ,𝑔 + log 10 𝑎 𝐶𝑙 -,𝑔 ) (I-80)

7. 3 . 3

 33 Electrode calibration Conditioning of electrodes Two combination electrodes with saturated potassium chloride as reference filling solutions are used in the present work. One of them was reconditioned with a refilled reference solution of 3 mol L -1 NaClO4 + 0.01 mol L -1 NaCl and applied to solutions of inert electrolyte being NaClO4. The other one was used in the inert solutions of NaCl. The conditioning of electrode was carried out in the following way: (a) The reference electrode was cleaned with distilled water, then with a solution of 0.01 mol L -1 NaCl and finally with the new reference solution in order to eliminate the K + present in the original reference solution (saturated KCl) which may cause the precipitation of KClO4 inside the liquid junction. (b) The freshly prepared reference solution was added into the reference electrode with several grains of AgCl. (c) The electrode was immersed in a solution of 3 mol L -1 NaClO4 + 0.01 mol L -1 NaCl slightly acidified to pH ≈ 2-3 during 24h.

  Fig. I -14 Calibration curves for (a) the reconditioned electrode used in NaClO4 and (b) the original electrode used in NaCl.

  temperature will change the buffer pH value due to the influence on the hydrogen ion activity. The temperature coefficient dpH/dt are shown in Table I -8, with which the pH values for each buffer are recalculated to different temperatures. It should be noted that the default temperature on the pHmeter should be reset before performing the new calibration at different temperature from 22°C. All the buffer solutions used in this work are valid within the experimental temperature range. The Fig. I -15 illustrates the calibration curves using the normal combination electrode (saturated KCl as reference filling solution). The slopes and intercepts are listed in Table I -7 at different temperatures.

Fig. I - 15

 15 Fig. I -15 Calibration lines in the temperature range of 5-50°C and the pH values corresponding to the zero point of electrode at 5°C and 50°C.

7. 3 . 4

 34 Measurements of liquid junction potential at different ionic strengthsThe liquid junction potentials for the reconditioned and the original glass electrodes have been measured with NaClO4/NaCl solutions of different ionic strengths. These solutions are specially prepared at the same [H + ] being 0.01 mol L -1 of ionic strength from 0.01 to 3 mol L -1 .The difference ∆Ej is calculated from the Eq. I-81 from which the measured potential of a specific ionic strength solution is compared to that of the solution of I = 3 mol L -1 . Due to the strong dependency of activity on ionic strength, pcH (= -log10[H + ]) is used instead of pH (= -log10a(H + )) throughout the experiments. The aqueous information and experimental measurements of Emes are reported in TableI -9. These measures allow establishing, for the reconditioned/original glass electrode with 3 mol L -1 NaClO4/KCl as electrode filling solutions: ∆Ej -rA0 log10(a2/a1) = 19.20 log10(3.50/Im) + 7.04/∆Ej -rA0log10(a2/a1) = 11.86 log10(3.19/Im) + 12.03, thus defining the function f(Isolution) with respect to ionic strength, as shown in Fig. I -16.

Fig

  Fig. I -16 Influence of ionic strength on the junction potential (a) with 3 mol L -1 NaClO4 as electrode filling solution in the reconditioned glass electrode and (b) with 3 mol L -1 KCl as electrode filling solution in the original glass electrode.

3. 1

 1 MaterialsA stock solution of natural U(VI) was prepared after dissolution of U3O8 in 37 % hydrochloric acid (Sigma-Aldrich, ACS reagent) during one week. The aqueous concentration of uranium is measured by inductively coupled plasma mass spectroscopy (ICP-MS). The total uranium concentration in the experiments is maintained at 50 μmol kgw -1 by diluting in NaCl solutions of different ionic strengths (0.1, 0.25, 0.5, 0.75 and 1 mol kgw -1 ), prepared with anhydrous NaCl (Sigma-Aldrich, ACS reagent, ≥ 99 %) and Millipore deionized water (Alpha-Q, 18.2 MΩ/cm). The observed spectral behaviour of U(VI), assumed to be the cumulative contribution of aqueous species, under saturated with respect to schoepite and calcite, is therefore an essential prerequisite for further acquisition. No precipitation effects has occurred at the present concentration of U(VI).The pH values of the sample solutions were controlled in the range 7.4-9.0 with sodium hydroxide (Analytical Grade, Fisher) and diluted hydrochloric acidthe pH measurement protocol are detailed afterwards. The desired pH value is regularly checked and maintained unvaried in each analysed system throughout the experiments. The last measurement of pH value is done before TRLS analysis. Carbonate is introduced by adding appropriate amounts of freshly prepared NaHCO3 solution (Analytical Grade, Fisher), calculated on the basis of equilibrium between aqueous solution and atmospheric CO2(g). Calcium chloride (Sigma-Aldrich, ACS reagent, ≥ 99 %) is used to control the concentration of calcium. Sample solution is prepared in 15 mL polyethylene bottle and titrated with adequate amount of CaCl2 solution.

  10 °C) is used to detect the time-resolved fluorescence signal. Emission spectra are recorded as a function of gate delay D (≥ 25 ns, minimum for preventing the Raman scatter peak of H2O) at the fixed gate width W = 1 μs. Following the expression of fluorescence signal in Eq. (II -1), the collected spectra are then integrated over the same wavelength span to deduce the fluorescence intensity at various delay time. The concentration of U(VI) is low enough to eliminate the absorption effect of excitation and emission wavelength, i.e. the pre-and postfilter effects. FI = ∑ FI 0 τ i exp (-

  in atmospherically-equilibrated aqueous solution, owing to the dissolution of CO2(g) in H2CO3(aq), HCO3 -, and CO3 2-. In contrast, in neutral to mildly alkaline conditions without calcium, complicated speciation of uranium(VI) driven by hydrolysis and oligomerization, even possible precipitation of UO3:2H2O(cr), would render the prediction of individual concentration for all the involved species difficult. For this reason, direct fluorescence measurement of uranyl-carbonate solutions prepared at investigated pH values without calcium is untenable, as characteristic fluorescence properties would be inevitably interfered by hydroxo uranyl species. Fluorescence decay spectra of a sample prepared at pH 7.6 were acquired over a large range of delay time, where the spectral profiles were significantly affected as shown in Fig. II -S2 of the SI.

  CO3)3(aq),NaCl). The activity coefficient is expressed as follows, log10 =zj 2 DH + (j,k,Im) mk (II -9) with DH the Debye-Hückel term,

  Five series of solutions are separately prepared at designed ionic strengths. The fluorescence spectra of aqueous Ca-UO2-CO3 complexes at Im = 0.5 mol kgw -1 NaCl are shown in Fig. II -1those of the four other series are shown in Fig. II -S3 of the SI. A total of 100 sample fluorescence spectra acquired at D = 25 ns were imported into Origin 9.0 software for Gaussian-based peak separation analysis. Characteristic peak wavelengths found at 464.31 ± 0.

Fig. II - 1

 1 Fig. II -1 Measured luminescence emission spectra of uranium(VI) at various calcium concentrations and pH values, P(CO2) = 10 -3.5 atm, and Im = 0.5 mol kgw -1 NaCl; initial delay time of D = 25 ns and gate width W = 1 μs, number of accumulation is 1000.

Fig. II - 2 .

 2 Fig. II -2 . Logarithmic ratio R in terms of luminescence intensity and Ringböm coefficient plotted as a function of log10[Ca 2+ ] (mol kgw -1 ) at investigated NaCl ionic strengths: [U]total = 50 µmol kgw -1 , P(CO2) = 10 -3.5 atm The respective decay-times are deduced by averaging the measured decay-times of correspondent sample solutions. It is assessed that in 0.1 mol kgw -1 NaCl medium, UO2(CO3)3 4-is showing a fast decay-time of (UO2(CO3)3 4-) =(11.04 ± 0.11) ns, compared with(CaUO2(CO3)3 2-) = (28.30 ± 0.57) ns and (Ca2UO2(CO3)3(aq)) = (46.7 ± 2.7) ns. The possible quenching effect of chloride on luminescence decay-time make it moderately meaningful to compare our results with those in literature, generally in non-complexing medium (0.1 mol L -1 Na/HClO4),6,8 albeit measured decay-times are numerically comparable with each other. In addition, a noticeable suppression of Ca-UO2-CO3 complexes decay-times is observed at larger NaCl concentrations that gives a clue to information such as quenching rate constant, which would require deeper attention in further works.

Fig. II - 3

 3 Fig. II -3 Mono-exponential fits of luminescence intensity decay of (a) UO2(CO3)3 4-at [U(VI)] = 50 µmol kgw -1 and pH = 9.0, and (b) Ca-UO2-CO3 complexes at [U(VI)] = 50 µmol kgw -1 , log10[Ca 2+ ] = -2.06, pH = 7.60 and Im = 0.1 mol kgw -1 NaCl. The insets show the waterfall plots of luminescence emission spectra as a function of delay time.

Fig. II - 4

 4 Fig. II -4 Luminescence emission spectra normalized to the same area at D = 25 ns of (Ca-)UO2-CO3 species at (a) Im = 0.2 mol kgw -1 , and (b) 1 mol kgw -1 ; Ca0, Ca1, and Ca2 stand for UO2(CO3)3 4-, CaUO2(CO3)3 2-, and Ca2UO2(CO3)3(aq), respectively.

Fig. II - 5

 5 Fig. II -5 Extrapolation to infinite dilution of the experimental data of (a) log10K1.1.3 + 16D and (b) log10K2.1.3 + 24D as a function of molal ionic strength. The predominance diagram at Im = 0.2 mol kgw -1 NaCl in Fig. II -6 is constructed from our newly determined log10°n.1.3 and  values using PhreePlot, 59,60 which includes PhreeqC 61,62 that handles the SIT. The apparent precipitation region of calcite, limited by the black dashed line, is imposed by the equilibrium between the atmospheric CO2(g) and the aqueous solutionthe four other series are shown in Fig. II -S5 of the SI. As expected, the increase of calcium concentration impedes the precipitation of UO3:2H2O(cr)schoepite. It was observed that the

- 1 ,

 1 illustrated by the repartition of the data in Fig. II -S6 of the SI. Our value for (CaUO2(CO3)3 2-,Na + ) = 0.29 ± 0.11 is higher than any other observed valuesand markedly higher than (UO2(CO3)2 2-,Na + ) = -(0.02 ± 0.09) -; only (Cu(OH)4 2-,Na + ) or (UO2F4 2-,Na + ) are contained in the standard deviation. Other multi-centred complexes as, e.g., UO2Edta 2-, are

  -S7(a)), and thermodynamic constants (Fig. II -S7(b)). It is clearly shown that our results

Fig. II - 7

 7 Fig. II -7 Variations of activity coefficients (a) and formation constants (b) for the MnUO2(CO3)3 (4-2n)-complexes (M being Mg or Ca) as a function of molal ionic strength.

  compositions are shown in Fig. II -8. The speciation results are markedly different because of the differences in total carbonate, calcium and magnesium concentrations, and Eh values. In the case of the Boom Clay water (Fig. II -8(a), (b)), 34 even with a lower Eh value, the carbonate concentration in the Boom Clay water stabilizes the different triscarbonatouranyl complexes at the expenses of U(OH)4(aq). The marked change in uranium speciation between CaUO2(CO3)3 2-and UO2(CO3)3 4-at pH 8.2 corresponds to the precipitation of dolomite (CaMg(CO3)2), which also affects the MgUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) complexes. In the case of the Callovo-Oxfordian water (Fig. II -8(c), (d)),35 the lower carbonate content but high calcium concentration yields to the predominance of MnUO2(CO3)3 (4-2n)-complexes at pH values lower than 8.5, and to the predominance of U(OH)4(aq) at higher pH values. In both cases, the MgUO2(CO3)3 2-complex should be minor; the SrUO2(CO3)3 2-complex should be even minor.

Fig. II - 8

 8 Fig. II -8 Solubility of UO2(am,hyd) (a,c) and aqueous speciation of uranium (b,d) between pH 7 and 9 for Boom 34 (a,b) and Callovo-Oxfordian 35(c,d) clay equilibrium waters using thermodynamic data from the Thermochime 9b database,80 implementing MgUO2(CO3)32-and SrUO2(CO3)3 2-,12 correcting the CanUO2(CO3)3(4-2n)-data with thermodynamic constants values from TableII-1, and implementing specific ion interaction coefficient from TableII-3: P(O2) = 10 -68.67 (a,b) and 10 -65.83 (c,d) atm.

(

  CaMg(CO3)2) and calcite (CaCO3) are showing positive saturation indices at pH values above 7.6 and 7.8, respectively (Fig. II -S7 of the SI). This means that the initial seawater composition is oversaturated with regards of these two phases. For this exercise, the precipitation of the carbonate phases of alkaline earth metals was not required.This exercise confirms the general predominance of CanUO2(CO3)3(4-2n)-in seawater and enlightens the need of comparable studies on the MgnUO2(CO3)3 (4-2n)-complexes at varying ionic strengths. As Mg/Ca concentration ratio is ca. 5 from Millero, et al. seawater composition, 81 and °( is ca. 12 according to the estimation from Dong and Brooks, 12 MgUO2(CO3)3 2-should be significantly present in seawatersee the theoretical speciation proposed in Lee, et al.7 Further works on the MgnUO2(CO3)3 (4-2n)-complexes vs. ionic strength are necessary to reduce the uncertainties on thermodynamic constants, and avoid the approximations for the specific ion interaction coefficients values.

Fig. II - 9

 9 Fig. II -9Theoretical speciation of uranium in a standard seawater composition, 81 P(CO2) = 10 -3.24 atm from the total carbonate concentration, and using the thermodynamic data from the Thermochimie 9b database,80 correcting the CanUO2(CO3)3(4-2n)-data with thermodynamic constants values (TableII -1) and implementing the specific ion interaction coefficients (TableII-3) proposed in this work.

  : the measured fluorescence emission spectra at atmospheric CO2(g), various Ca concentration and pH values, and 5 different ionic strength values; the predominance plots of the Ca-UO2-CO3 system under the experimental conditions; frequency distribution of the 𝜀(X 2-, Na + ) values from literature compared to the 𝜀(CaUO 2 (CO 3 ) 3 2-, Na + ) determined in this work; and evolution of the saturation indices of calcite (CaCO 3 ) , nesquehonite (MgCO 3 : 3H 2 O) , and dolomite (CaMg(CO 3 ) 2 ) from the theoretical speciation in Fig. II -9 of the main text; the stepwise formation constants values from Fig. II -5 of the main text; the ε(X 2-,Na + ) values from literature.

Fig

  Fig. II -S1 Species distribution at [U(VI)] = 50 μmol kgw -1 in preliminary calculations with stability constants for UO2(CO3)3 4-, UO2(CO3)2 2-, UO2(CO3)3(aq), and (UO2)2CO3(OH)3 -listed in TableII-1, and those for CanUO2(CO3)34-2n taken from the work of Lee and Yun.1 Other aqueous information is in Table II -S1.

Fig

  Fig. II -S2 (a) Fluorescence decay spectra of the sample at [U(VI)] = 50 μmol kgw -1 , pH = 7.6 equilibrated with atmospheric CO2(g), recorded with λex = 450 nm from D = 20 ns to D = 210 ns. (b) Red line: spectrum recorded at D = 20 ns. Blue line: spectrum at D = 210 ns. Green line: subtraction of the blue line from the red one, also characteristic spectrum of UO2-CO3 species.

Fig

  Fig. II -S4 Hypsochromic-shifted fluorescence spectrum of UO2(CO3)3 4-([U(VI)] = 50 μmol kgw -1 , pH = 9, Im = 0.1 mol kgw -1 NaCl) compared with that of UO2 2+ ([U(VI)] = 0.1 mmol kgw -1 , pH = 1) in HClO4.

  2- and Ca2UO2(CO3)3(aq), were determined in NaClO4 medium at various ionic strengths using time-resolved laser-induced luminescence spectroscopy (TRLS)also known as time-resolved laser-induced fluorescence spectroscopy (TRLFS). As in a previous study, the potential precipitation of schoepite (UO3•2H2O) and calcite (CaCO3) were avoided via titration of triscarbonatouranyl complex by Ca 2+ at varying pH values. The Ringböm coefficients relative to UO2(CO3)3 4-were individually evaluated for test sample conditions. Steadily enhanced luminescence intensity and increased decay-times were representative of complexation processes. The stoichiometry of calcium was quantified by slope analysis, the measured intensity of which was corrected by the corresponding Ringböm coefficient. The conditional formation constants, i.e. log10 Kn.1.3, were then assessed after rounding off the slope values to their nearest integers. Determinations of cumulative formation constants at infinite dilution log10 °n.1.3, and specific ion interaction parameters ε were based on the experimental origin and slope values, respectively, over the range 0.1-2.46 mol kgw -1 NaClO4 using the specific ion (CaUO2(CO3)3 2-) = 27.26 ± 0.04 and log10 °(Ca2UO2(CO3)3(aq)) = 30.53 ± 0.06. The specific ion interaction coefficients are estimated to be ε(CaUO2(CO3)3 2-,Na + ) = (0.02 ± 0.04) kgw mol -1 and ε(Ca2UO2(CO3)3(aq),NaClO4) = (0.18 ± 0.07) kgw mol -1 . These latter values are different from the ones that were our previously obtained in NaCl, and underlying causes are discussed from different aspects. This work provides valuable information to address the interaction effects between Ca-UO2-CO3 species and 1:1 type electrolytes. It is suggested that the affinity of the cation in background electrolyte with CanUO2(CO3)3 (4-2n)-(n = {1;2}) are to be taken into consideration at high ionic strengths, especially for the globally non-charged species.

(

  ACS reagent) to prepare the stock solution of U(VI). A brilliant greenish-yellow U(VI) liquid solution was obtained at the end of 48 hours. Inductively coupled plasma mass spectroscopy (ICP-MS) was used to measure the aqueous concentration of uranium after multiple dilutions of the concentrated U(VI) solution. Seven series of NaClO4 solutions of different ionic strengths (i.e. 0.1, 0.51, 1.05, 1.62, 2.21, 2.46, and 3.5 m) were prepared by dissolving anhydrous NaClO4

  coefficients and stability constants by linear regression analysis. The results at the highest ionic strength of this work, i.e. 3.5 m NaClO4, were excluded due to the lack of the second slope of two in slope analysisvide post and Fig. III -S1 in the SI. As shown in Fig. III -1, the emission spectra at D = 25 ns illustrate the successive complexations of calcium ions with triscarbonatouranyl moiety. The expected spectral and temporal characteristicsbetter resolved spectrum, continuously enhanced luminescence intensity and slower decay-times -, agree well with the changes of stoichiometry from monocalcium-to dicalciumtriscarbonatouranyl complexes in aqueous solution.

Fig. III - 1

 1 Fig. III -1 Evolution of luminescence emission spectra during complexation of Ca(II) with UO2(CO3)3 4-at different ionic strengths: the aqueous conditions where the ternary species dominate are listed inTable III -3.

Fig. III - 2

 2 Fig. III -2 Plot of logarithmic ratio R as a function of log10[Ca 2+ ] (m) at (a) Im = 0.10 m NaClO4 (b) Im = 0.51 m NaClO4 (c) Im = 1.05 m NaClO4 (d) Im = 1.62 m NaClO4 (e) Im = 2.21 m NaClO4 (e) Im = 2.46 m NaClO4 : [U]total = 50 μm, P(CO2) = 10 -3.5 atm.

5 m = 1 .

 51 37 in NaClO4 compared to Δlog10 K 0.1 m = 2.43 in NaCl. With increasing electrolyte concentration, the difference becomes evident: Δlog10 K 0.91 in NaClO4 compared to Δlog10 K 0.5 m = 1.79 in NaCl, and Δlog10 K 1.0 m = 1.81 in NaClO4 compared to Δlog10 K 1.0 m = 1.51 in NaCl. This tendency indicates that the association between Ca 2+ and CaUO2(CO3)3 2-is more likely to occur in NaClO4 than in NaCl, since the values of Δlog10 K are more important in NaClO4. The extrapolation of the cumulative constants to infinite dilution is shown in Fig. III -3(a,b) with the variation of log10 1.1.3 + 16DH and log10 2.1.3 + 24DH vs. ionic strength in NaClO4 medium (red circles). The cumulative formation constants, i.e. log10 1.1.3 and log10 2.1.3, were calculated by adding log10 0.1.3 from Guillaumont, et al. 29 to either log10 K1.1.3 or log10 K2.1.3, with uncertainties determined by propagation of errorthe values of log10 0.1.3 at different Im were calculated using the SIT approach and listed in Table III -4. The linear relationships enable the evaluations of log10 ° (intercept, Im = 0) and ∆ε (slope) within the SIT framework. Systematically, propagation of errors is applied to assess the uncertainty range from infinite dilution up to 3 m NaClO4. The resulting stability constants are log10 °1.1.3 = 27.26 ± 0.08 and log10 °2.1.3 = 30.53 ± 0.12; the Gibbs energies of reaction are ΔrG°m(CaUO2(CO3)3 2-) = -(155.601 ± 0.457) kJ mol -1 and ΔrG°m(Ca2UO2(CO3)3(aq)) = -(174.267 ± 0.685) kJ mol -1 ; and the Gibbs energies of formation are ΔfG°m(CaUO2(CO3)3 2-) = -(3 244.658 ± 2.394) kJ mol -1 and ΔfG°m(CaUO2(CO3)3(aq)) = -(3 816.130 ± 3.050) kJ mol -1 . These values are in excellent agreement with our previous work in NaCl 8see Table III-5 and blue squares in Fig. III -3(a,b) -, and with other literature data within uncertainty limits.2,3,6,9,11

Fig. III - 3

 3 Fig. III -3 Specific ion interaction theory extrapolation to infinite dilution of the experimental data of (a,c) log10 1.1.3 + 16DH and (b,d) log10 2.1.3 + 24DH, for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), respectively, as a function of molal ionic strength for different ionic media: NaCl data (blue squares) are from Shang and Reiller 8 ; (a,b) extrapolation in individual ionic media (red line, present work; blue line from Shang and Reiller 8 ); (c,d) multilinear regression of NaClO4 and NaCl 8 data; dashed line are the propagation of errors from infinite dilution to 3 m, using the different regressions.

1 . 3 +

 13 16DH and log10 2.1.3 + 24DH vs. Im in NaCl 8from our previous work using the same experimental methodand NaClO4 in Fig. III -3 shows that the log10 °n.1.3 values are so similar that a weighted average would give log10 °1.1.3 = 27.20 ± 0.01 and log10 °2.1.3 = 30.49 ± 0.01. A multilinear regressionas in Hummel, et al. 62is leading 123 to the plots in Fig. III -3 (c,d) that give common log10 °n.1.3 values for NaCl and NaClO4, and the ε values for the ionic media. The common values of log10 °1.1.3 = 27.22 ± 0.04 and log10 °2.1.3 = 30.51 ± 0.06 can be proposed.

  from test samples prepared at Im = 0.1 m NaCl/NaClO4 and 1 m NaCl/1.05 m NaClO4. As shown in Fig. III -4, the normalized spectra display mostly the same profile in the entire acquisition range when the composition of aqueous solutions are nearly identical. This observation indicates that the principle spectral parameters are mainly determined by the proportion of aqueous binary/ternary species and not affected by the type or the concentration of solvent anion at least at the resolution of the spectroscopic set-up.

Fig. III - 4

 4 Fig. III -4 Luminescence emission spectra normalized to the same area at D = 25 ns, W = 1 μs of binary and ternary species at (a) Im = 0.1 m, and (b) Im = 1 m. Ca0, Ca1 and Ca2 stand for UO2(CO3)3 4-, CaUO2(CO3)3 2-, Ca2UO2(CO3)3(aq), respectively.

Fig. III - 5

 5 Fig. III -5 shows normalized spectra of chosen samples of equal calcium concentration (ca.6 × 10 -3 m) at increasing ionic strengths from 0.1 to 2.46 m NaClO4. The Gaussian-peak shape was used to fit the emission bands with Levenberg-Marquardt method and to perform the spectrum decomposition. As expected, the increase of Im hinders the complexation of calcium and results in a decrease of the proportion of aqueous calcium(II)triscarbonatouranyl(VI)

Fig. III - 5

 5 Fig. III -5 Luminescence emission spectra normalized to the same area at D = 25 ns of Ca-UO2-CO3 species: the aqueous compositions are listed in Table III -6; (a) the sample solutions having the similar composition are specifically selected at different ionic strengths; (b) the Ca(II) concentration is fixed at 6 ×10 -3 m; the variation of spectral feature vs. ionic strength was shown in Fig. III -6.

Fig. III - 6

 6 Fig. III -6 Maximum heights of the peaks of 10→00 (484 nm) and 10→01 (503 nm) transitions are obtained from the spectra in Fig. 5 (b) using Gaussian decomposition: Ca0, Ca1+Ca2 stand for UO2(CO3)3 4-, the sum of CaUO2(CO3)3 2-, Ca2UO2(CO3)3(aq), respectively.

  2-and Ca2UO2(CO3)3(aq) are respectively assessed by averaging the measured decay-times of the sample solutions giving the slopes of 1 and 2 (Fig. III -2). Whereas the decay-time of UO2(CO3)3 4-is evaluated by measuring the sample prepared at [U(VI)] = 50 μm, pH = 9 and P(CO2) =10 -3.5 atm at each ionic strength. The exact values of the three species are listed in Table III -S4. Contrary to the series of NaCl medium, 8 where the intensity decay curves were fitted by mono-exponential function, two exponential functions containing two sets of parameters, i.e. intensities FI1, FI2 and decay-times τ1, τ2, fit much better the sets of experimental data in NaClO4 series. A typical example of the decay-time analysis is presented in Fig. III -7. The more parametric bi-exponential function yields two decay-times of which the fast decay one (5.68 ns) has a very high initial fluorescence intensity (1.19 x 10 9counts) governing the total FI in short delay times (D < 25 ns) and the slow decay one have similar results as the mono-exponential fitting. It was verified a posteriori that in all bi-exponential fitting cases, one of the two components always have a much shorter decay-time and a higher initial intensity compared to the other one. However, the contribution of the fast decay component is nearly negligible (below 10%) and is only due to one data point at low delay after the laser flash -D = 25 ns. Thus, this fast-decay component is possibly an artefact in the acquisition process and should not be considered as a meaningful species in the decaytime analysis. The estimation for decay-times of the studied species is consequently based on single exponential decay fit in this study.

Fig. III - 7

 7 Fig. III -7 Mono-exponential and bi-exponential fits of luminescence intensity decay of Ca-UO2-CO3 complexes at [U(VI)] = 50 μm, log10[Ca 2+ ] = -1.88, pH = 7.48 and Im = 1.05 m NaClO4.

Fig. III - 8

 8 Fig. III -8 Averaged decay-times of sample solutions giving the slopes of 1 (red circles) and 2 (blue triangles) at investigated NaClO4 ionic strengths; the values for UO2(CO3)3 4-complex are decay-times of the sample prepared at [U(VI)] = 50 μm, pH = 9 and P(CO2) =10 -3.5 atm at each Im; the error bars of UO2(CO3)3 4-are too small to be shown; Exact values are listed in Table III -S4 of the SI.

4. 4

 4 Fig. III -S1 of the SI shows the slope analysis at Im = 3.50 m. Although the calcium concentrations are considered sufficient for the second complexation to occur, only the slope of 1 is obtained using the least-squares adjustment to our data. In the predominance plot (Fig. III -S2 of the SI), the test samples (red square) at the end of complexation could be affected by the onset of calcite precipitation, induced by the practical failure of fixing the pH values overly close to the steep plunge of pH near the neutral zone. One can note that 17 sample solutions have primarily satisfied the formation conditions of Ca2UO2(CO3)3(aq) in terms of calcium concentration and pH value. A striking fact is that the 13 experimental points giving the slope of one are all found set in the predominance region of Ca2UO2(CO3)3(aq). In addition, the last three points in slope analysis seem to reach a plateau in terms of intensity (Fig S1 in the SI),while the decay-times increased with the addition of Ca 2+ . For this reason, the intercept value obtained at Im = 3.50 m was not taken in the evaluation of log10 ° and ε. Nevertheless, the resulting constant would be included as an indicative point in the fits using different predictive models in an extended range of Im in the next section.

  developed for SIT ionic strength correction. Fig. III -9 depicts the distribution of experimental points on the [Ca]-pH plan of each series. The equilibrium between the atmospheric CO2(g) (P(CO2) = 10 -3.5 atm) and the aqueous solution is calculated to identify the precipitation region of calcite, imposed by the black dashed line. It seems evident that the increase of ionic strength enlarges the predominance domain of UO2(CO3)3 4-and narrows those of the ternary and polynuclear complexes -CaUO2(CO3)3 2-, Ca2UO2(CO3)3(aq), and (UO2)2CO3(OH)3 -illustrating the hindrance of ionic strength to the formation the calcium(II)triscarbonatouranyl(VI) complex. The domain limits are reflecting the experimental results. Nevertheless, it should be noted that at relatively high electrolyte concentrations (from 1.05 to 2.46 m NaClO4), the sample solutions giving the desired slopes cross the boundaries to some extent at the beginning of complexation.

Fig. III - 9

 9 Fig. III -9 Predominance plots obtained using PHREEPLOT 80 of Ca-UO2-CO3 system at [U(VI)] = 50 μm, P(CO2) = 10 -3.5 atm at ionic strength from 0.1 to 2.46 m NaClO4: experimental points giving slopes of 1 and 2 are highlighted with red squares and blue triangles, respectively; the black circles represent the beginning of titration where the binary complex dominates; dashed lines represents the line of predominance between Ca 2+ and calcite.

3

 3 and constant ε to deduce log10 n.1.3 at specific ionic strengths could be questioned for high Im situations. Several modified models based on the SIT approach are applied to better represent our experimental data and examine the dependence of interaction coefficients on ionic strength. The fit results of the one-parameter SIT model are compared with those of two-parameter models in Fig. III -10. Considering the significant suppression in terms of fluorescence intensity at Im = 3.50 m, only the conditional formation constants determined in the range of 0.1 to 2.46 m are fitted to the models. The relationships used to fit the data set are expressed as follows:

Fig. III - 10

 10 Fig. III -10 Fits of experimental data (a) log10 113 +16DH and (b) log10 213 +24DH by different models and extrapolation to 6 m.

  This supporting information contains four Tables: the liquid junction potentials measured for NaClO4 solutions of different ionic strengths (0.01 m ≤ Im ≤ 3.5 m) using the reconditioned electrode; experimental samples information -pH value, calculated Ringböm coefficients α, [Ca 2+ ] (mol•kgw -1 ) and deduced F0; information for samples in Fig. III -4; Decay-times of CanUO2(CO3)3 (4-2n)-determined at each Im ; and two Figures: the slope analysis results of Im = 3.5 m; the predominance plot of the 3.5 m experiment.

Fig

  Fig. III -S1 Plot of logarithmic ratio R as a function of log10[Ca 2+ ] (mol kgw -1 ) at Im = 3.50 m of NaClO4 medium: [U]total = 50 μmol kgw -1 , P(CO2) = 10 -3.5 atm.

3 = 26 .

 326 luminescence quenching effects generated by high concentrations of Cl -and ClO4 -, already observed during our earlier studies on CanUO2(CO3)3 (4-2n)-(Shang & Reiller, Dalton Trans. 49, 466; Shang et al., ibid 49, 15443). The cumulative formation constant determined is log10β°1.1.40 ± 0.07, and the specific ion interaction coefficients ε(MgUO2(CO3)3 2-, Na + ) = 0.19 ± 0.11 kgw mol -1 in NaClO4 and ε(MgUO2(CO3)3 2-, Na + ) = 0.09 ± 0.16 kgw mol -1 in NaCl. The 300 and 1800 lines mm -1 gratings were both used to acquire MgUO2(CO3)3 2-luminescence spectra, where the high-resolution 1800 lines mm -1 grating detected slight spectral shifts for the principal luminescent bands relative to CanUO2(CO3)3 (4-2n)-. The application of the consistent set of thermodynamic constants and ε values for MnUO2(CO3)3 (4-2n)-(M = Mg, Ca) are examined in different geochemical contexts, where Mg over Ca concentration ratio varies to help defining the relative importance of these species.

  . IV -S1 of the supplementary information (SI) -, the stability constant of MgUO2(CO3)3 2-also presents a marked anomaly in comparison with the log10βº (MUO2(CO3)3 2-) for the other stable alkaline earth cations -M = Ca, Sr, and Ba in Dong and Brooks14 in Fig. IV -S2 of the SI. The question raised here is whether the hydration structural nuances of Mg 2+ and Ca 2+ will be reflected in the formations of MnUO2(CO3)3(4-2n)-at high ionic strengths. The comparison of log10β° and ε for similar types of MnUO2(CO3)3 (4-2n)-complexes determined in different media allow achieving a clearer view of ion interaction mechanism, and help rationalizing the differences in interaction coefficients.

  intensities FI0 at D = 0 and decay-times τ for individual sample solution were obtained by integrating time-resolved luminescence spectra with the trapezoidal rule then fitting the resulting intensities FI at various delay times to the exponential decay functions. Fig. IV -1 shows the graphical relationship between log10R and log10[Mg 2+ ] at varying ionic strengths.

Fig. IV - 1

 1 Fig. IV -1 Plot of logarithmic ratio R as a function of log10[Mg 2+ ] (mol kgw -1 ) at (a) -(e) Im = 0.10 -1.0 mol kgw -1 of NaCl, (f) -(j) Im = 0.10 -2.21 mol kgw -1 of NaClO4: [U]total = 50 μmol kgw -1 , P(CO2) = 10 -3.5 atm.

  4 lists the total magnesium concentration and aqueous pH values where the linear relationship is satisfied. The slope values of the fitting results in Fig. IV -1 were rounded-off to unity, and the associated intercepts were assigned to the formation constants of MgUO2(CO3)3 2-at different ionic strengths in NaClO4 and NaCl. Table IV -5 summarizes the apparent successive constants log10K1.1.3, the cumulative constants log10β1.1.3, and the values of log10β1.1.3 + 16DH for the evaluations of log10β° and ε(MgUO2(CO3)3 2-,Na + ). The dependence of the conditional equilibrium constants on ionic strength is clearly shown. As Im increases, the interaction between Mg 2+ and UO2(CO3)3 4-is inevitably weakened as a result of the attraction from the oppositely charged ions in solution, i.e. Cl -/ClO4 -and Na + . This phenomenon can be numerically interpreted by the decreasing trend of log10K1.1.3 vs. Im, as shown in Table IV -5.

Fig. IV - 2

 2 Fig. IV -2(a) shows the variation of log10β1.1.3 + 16DH as a function of Im in NaCl and NaClO4, where the values of log10β1.1.3 is the sum of log10β0.1.3 and log10K1.1.3 and DH are recalculated with respect to the specific ionic strengths. Error propagation was used to evaluate the uncertainty ranges. The obtained equilibrium constants are determined as log10β°1.1.3 = 26.27 ± 0.06 in NaClO4 and log10β°1.1.3 = 26.57 ± 0.01 in NaCl. The specific ion interaction

Fig. IV - 2

 2 Fig. IV -2 Extrapolation to infinite dilution of the experimental data of log10β1.1.3+ 16DH as a function of ionic strength (a) by individual regression in NaCl and NaClO4 and (b) by multilinear regression.

  stated that MgUO2(CO3)3 2-and Mg2UO2(CO3)3(aq) are supposed to show almost indistinguishable luminescence spectra. The steric hindrance effects of Mg 2+ can be reflected by comparing the stability constants with various ligands with other elements in the alkaline earth group. The log10β° values often present a regular sequence from Ca to Ba with an anomaly occurred for Mg (see Fig. IV -S1 of the SI).

Fig. IV - 3

 3 Fig. IV -3 illustrates the complexation of Mg 2+ with UO2(CO3)3 4-in spectral domain using the 300 lines mm -1 grating. The progressively enhanced luminescence intensities, betterresolved spectra and steadily narrowing bands evidenced the binding process of alkaline earth metal ions. The spectral evolution pattern for Mg-UO2-CO3 complexes are in line with previous observations on alkaline earth cations. 2,9,10,12,24,25 Gaussian-based peak separation analysisimplemented in Origin 9.0 softwarewas applied to the luminescence spectra at D = 25 ns.

Fig. IV - 3

 3 Fig. IV -3 Graphical representations of luminescence emission spectra on complexation of Mg(II) with UO2(CO3)3 4-at investigated ionic strengths: (a) -(e) Im = 0.10 -1.0 mol kgw -1 of NaCl, (f) -(j) Im = 0.10 -2.21 mol kgw -1 of NaClO4, acquired at D = 25 ns and gate width W = 1μs, number of accumulation is 1000. The aqueous conditions where the MgUO2(CO3)3 2-complex dominate were listed in Table IV-4.

Fig. IV - 4

 4 Fig. IV -4 Gaussian decomposition of luminescence spectrum of the sample: U(VI) = 50 μmol kgw -1 , Mg(II) = 0.00149 mol kgw -1 , P(CO2) = 10 -3.5 atm, pH = 8.05 at 0.1 mol kgw -1 NaCl, acquired at delay time D = 25 ns and gate width W = 1μs, number of accumulation is 1000. The decomposition results were listed in Table IV -S3 of the SI.

Fig. IV - 6

 6

Fig. IV - 6

 6 Fig. IV -6 Variations of decay-times obtained by averaging the sample solutions giving the slopes of 1 at investigated NaCl and NaClO4 ionic strengths.

  3 and ε values of MnUO2(CO3)3(4-2n)-(M = Mg, Ca, n = {1;2}) determined in NaCl from our previous work for Ca,24 and those for Mg in the present work. The origin of data is detailed on the web pages (http://www.thermochimie-tdb.com). It is worth noting that the originally implemented stability constants of CanUO2(CO3)3 (4-2n)-(n = {1; 2}) in ThermoChimie 10a database were taken from Dong and Brooks,14 whilst the MgUO2(CO3)3 2-constant from the same authors was not implemented; no ε values were included for the MnUO2(CO3)3 (4-2n)-complexes, and were thus considered as nil in calculations.As illustrated in Fig.IV -7, the relative proportion of MgUO2(CO3)3 2-slightly increased ca. 0.5 unit (Fig.IV -7 (a, c)) (7-8% in percentage, see Fig.IV -7 (b, d)) in both contexts, compared with the results in our previous worksee Fig.8of Shang and Reiller.24 However, general distribution pattern did not exhibit strong alternation of the concerned species. As noted in Reiller and Descostes,7 the high concentration of carbonateor activity of either HCO3 -or CO3 2-in Boom Clay context assures the predominance of MnUO2(CO3)3 (4-2n)-complexes (n = {0;1;2}) throughout the pH values, even with a lower redox potential. Whereas in Callovo-Oxfordian context the higher Ca and lower carbonate concentrationsor Ca 2+ , and either HCO3 - or CO3 2-activitiesassure the predominance of Ca2UO2(CO3)3(aq) complex at pH values lower than 8.5, and then of U(OH)4(aq) in higher basicity because of a lower carbonate content. It can be seen in both cases that the proportions of MgUO2(CO3)3 2-in solution is significantly higher than in our previous exercise. This is more evident for the Boom clay context compared to the Callovo-Oxfordian context where the ratio of Ca/Mg is highersee water compositions in SI in Reiller and Descostes.7

Fig. IV - 7

 7 Fig. IV -7 Solubility of UO2(am,hyd) (a,c) and aqueous speciation of uranium (b,d) between pH 7 and 9 for Boom (a,b) and Callovo-Oxfordian (c,d) clay equilibrium waters using thermodynamic data from the Thermochime 10a database, implementing SrUO2(CO3)3 2-,14 correcting the MgUO2(CO3)3 2-and CanUO2(CO3)3(4-2n)-data with thermodynamic constants values from determined in this work and our earlier work,24 and implementing specific ion interaction coefficient from TableIV-3: (a,b) P(O2) = 10 -68.67 atm and (c,d) 10 -65.83 atmsee Shang and Reiller 24 for details.

Fig. IV - 8

 8 Fig. IV -8 represents the distribution of uranium species as a function of pH in seawater. The MgUO2(CO3)3 2-and CaUO2(CO3)3 2-complexes share nearly equal proportions of 35% at pH values from 7.5 to 9, while the amounts of UO2(CO3)3 4-and Ca2UO2(CO3)3(aq) represent less than one-third in total U(VI). The slight decrease of CanUO2(CO3)3 (4-2n)-at pH > 9 was resulted from the decrease of [Ca 2+ ] in solution as a consequence of dolomite and calcite formations at high pH conditions. Henceforth, It seems that the speciation of uranium in seawater is strongly influenced by both MgUO2(CO3)3 2-and CaUO2(CO3)3 2-complexes.

Fig. IV - 8

 8 Fig. IV -8Theoretical speciation of uranium in a standard seawater composition, 38 P(CO2) = 10 -3.5 atm from the total carbonate concentration, and using the thermodynamic data from the ThermoChimie 10a database, correcting the MgUO2(CO3)3 2-, CanUO2(CO3)3(4-2n)-data with thermodynamic constants determined in this work (Mg) and our earlier work.24,25

  9 and Lee, et al.10 listed in the Table IV -S7 of the SI. As shown in Fig. IV -S3 of the SI, the (M-)UO2-CO3 species differ considerably in distributions. The CaUO2(CO3)3 2-complex dominates the aqueous uranium distribution over 90% at pH values above 7.3 at the expenses of other triscarbonatouranyl species using the constants from Dong and Brooks 14,15 (see Fig. IV -S3(b) of the SI), while its proportion decreased to 62% using the constants from Lee and Yun 9 and Lee, et al. 10 (see Fig. IV -S3(a)

  description of activity correction, which was clearly not appropriate, but inevitable in view of the U(VI)-NOM model, and software used. They proposed that the distribution of uranium species was majorly dominated by the ternary complexes MnUO2(CO3)3(4-2n)-(M = Mg, Ca, n = {0; 1; 2}) in the collected samples. For this reason, the U(VI)-NOM complexes are not considered in this exercise.

  Fig. IV -9, the distribution of U(VI) are showing similar patterns but with different amplitudes

Fig. IV - 9

 9 Fig. IV -9 Evolution of repartitions of uranium species in selectedCanadian aquitard waters at 5°C39,40 as a function of depth: (a) using the ThermoChimie 10a database implemented with the log10β°n.1.3 determined in this work; and (b) using the PRODATA 1.1.0.4 database.7 The values for ∆rHm were listed in Table IV -S8 of the SI.

Fig. IV - 10

 10 Fig. IV -10 compares the species percentages for the selected water samples using different constants for log10β°(MnUO2(CO3)3 (4-2n)-) (M = {Mg 2+ , Ca 2+ } , n = {1;2} ). The higher log10β°(MgUO2(CO3)3 2-) determined in this work than the one used by Coyte, et al.41 leads to a more abundant MgUO2(CO3)3 2-and less abundant CanUO2(CO3)3(4-2n)-in modelling results.

Fig. IV - 10

 10 Fig. IV -10 Evolution of the repartitions of uranium species in selected Indian groundwater calculated (a) using the llnl.dat in the work of Coyte, et al. 41 (b) using the Thermochimie 10a database with the log10β°n.1.3 determined in this work and our earlier work.

CaUO2(CO3)3 2 -Fig

 2 Fig. IV -8 log10β° 27.20 a 30.49 a 26.57 p.w. / ε 0.29 ± 0.11 a 0.66 ± 0.12 a 0.31 ± 0.02 p.w. / Fig. IV -S4(a) log10β° 27.27 b 29.81 b 25.8 c 27.1 c ε -0.02 b 0 b -0.02 c 0 c Fig. IV -S4(b) log10β° 27.18 d 30.70 d 25.02 e / ε -3 d 0 d -3 e / a Shang and Reiller; 16 b Lee and Yun; 18 c Lee, et al.; 19 d Dong and Brooks; 15 e Dong and Brooks 20 Table IV -S8 Thermodynamic constants used in theoretical speciation of Canadian aquitard case in Fig. IV -9 of the main text and in Fig. IV -S6. CaUO2(CO3)3 2- Ca2UO2(CO3)3(aq) MgUO2(CO3)3 2-

Fig

  Fig. IV -S1 Linear relationship between the formation constants log10K° for Mg and Ca with different ligands. The values of log10Kº are listed inTable IV -S4.

Fig

  Fig. IV -S2 Stability constants of alkali earth metal ions and different ligands -F -, OH -, SO4 2- , CO3 2-as a function of 1/R with R being the size of the ionic radius. The values of log10Kº are listed in Table IV -S5.

Fig

  Fig. IV -S3 Stability constants of alkaline earth metal ions and UO2(CO3)3 4-as a function of 1/R with R being the size of the ionic radius. The values of log10Kº are listed inTable IV -S5.

Fig

  Fig. IV -S6 Predominance plots of uranium species at the fixed values of log 10 {a H + a HCO 3 -} corresponding to the aqueous conditions of the collected samples in Canadian aquitard site. Thermochimie 10a database file29 and SIT ionic strength correction were used in calculations with implemented thermodynamic constants of MgUO2(CO3)3 2-, CanUO2(CO3)3(4-2n)-listed in Table IV -S8. The boundaries plotted in dashed lines represent the simulation results at 5°C and the solid lines at 25°C. The ionic strength varies from 1.38 mol kgw -1 at 29 m to 0.06 mol kgw -1 to at 2.3 m.

  Before spectrum acquisition, the cuvette was placed in the thermostated sample holder during at least 20 min. The temperature was maintained at a constant value by circulating water through the thermostat (Peter Huber Kältemaschinenbau AG, Germany). The sample cuvette was thermally equilibrated with the circulating water throughout the measurements for minimizing the temperature fluctuation on uranyl fluorescence. The laser excitation source in this study is a frequency-tripled Nd:YAG laser at 355 nm (Surelite, Continuum, USA) providing a 5 ns pulse at 10 Hz of about 170 mJ as energy. An optical parametric oscillator system (Horizon, Continuum, USA) was used to tune the excitation wavelength at λexc = 450 nm. The fluorescence beam energyaverage of 100 laser shotswas read on a Joule-meter RJP-734 (Laser Probe, Inc.). The 300 lines per mm grating was selected in monochromator spectrometer (Acton). The resulting measured fluorescence signal, recorded at a delay time D = 25 ns with the gate width W = 1 μs, was detected by an intensified CCD camera (Andor, UK) maintained at -10°C by Peltier effect.The recorded spectra were integrated over the emission wavelength range at varying D values, with the trapezoidal method implemented in Origin 9.0 software (OriginLab Corporation, USA), to evaluate the fluorescence intensity at various delay times. The leastsquare fitting process (Levenberg-Marquardt algorithm) was used to calculate the luminescence intensity at D = 0 and decay-time τ by fitting the intensity results to the exponential fluorescence decay function, expressed as already discussed,[34][35][36] FI = ∑ FI 0 τ i exp (-

3 °+

 3 Values of the Debye-Hückel parameters A(T) at different temperatures are shown in Table V -S6 of the SI. The relationship between conditional formation constants and ionic strength is expressed by log 10 K n.1.3 = log 10 K n.1.Δz 2 D H -ΔεI m (V -7)

  four models were employed to describe the temperature-dependent log10K°n.1.3 measured in the temperature range 10-50°C for CanUO2(CO3)3 (4-2n)-and 5-30°C MgUO2(CO3)3 2-and used to cautiously extend their formation constants at higher temperatures. The simplified Van't Hoff equation, assuming constant ΔrH°m with temperature and zero heat capacity, is shown in Eq. V -10. log 10 K°(T)= log 10 K°(T 0 ) -∆ r H m °(T 0 )

Fig. V - 1

 1 Fig. V -1 and Fig. V -2 illustrate the evolutions of luminescent emission spectra of MnUO2(CO3)3 (4-2n)-complexes. The formation constants for CaUO2(CO3)3 2-were measured in temperature range of 10-50°C, while those for MgUO2(CO3)3 2-were achieved in a relatively narrower temperature range (5-30°C), because of a pronounced decrease in the luminescence emission intensity, and a strong deformation of characteristic spectra at T > 30°C that caused difficulties in interpreting the luminescence evolution.

Fig. V - 1

 1 Fig. V -1 Evolution of luminescence emission spectra during complexation of Ca(II) with UO2(CO3)3 4-at different temperature (a)-(f): 10-50°C.

Fig. V - 3

 3 Fig. V -3 Luminescence emission spectra at higher temperatures than the ambient of the sample: [U(VI)] = 50 μm, [Ca 2+ ] = 6.13 10 -4 mol kgw -1 , pH ≈ 8 with the dominant species CaUO2(CO3)3 2-in solution.

Fig. V - 4

 4 Fig. V -4 Luminescence emission spectra of UO2(CO3)3 4-normalized to the same area at D = 25 ns; the concentration of [U(VI)] increased at high temperatures (40-50°C) to enhance the emission intensity.

2 -

 2 Fig. V -5 Pure emission spectra obtained from MCR-ALS analysis of the complexation of Ca 2+with UO2(CO3)3 4-at 50°C.

4. 2

 2 Evaluation of equilibrium constants at different temperaturesThe graphical representations of linear regression results at different temperatures are shown in Fig.V -6for CanUO2(CO3)3(4-2n)-and in Fig.V -7for MgUO2(CO3) 2-. The stability constants at 22°C of MnUO2(CO3)3 (4-2n)-had been measured during our previous studies.34,36 The 2σ error bars are too small to be presented due to the trapezoidal integration method used to deduce the value of FI0 (the fluorescence intensity at D = 0). Linear fits were achieved by weighted leastsquares regression analysis with estimated parameters of slope and intercept mean values with 95% uncertainties. After rounding off the slope values to the integers of 1 and 2, i.e. the stoichiometry of coordinated Mg 2+ /Ca 2+ , we obtained the conditional formation constants of MnUO2(CO3)3 (4-2n)-listed in Table V -1 with the log10K°n.1.3 extrapolated to infinite dilution by SIT.

Fig. V - 6

 6 Fig. V -6 Plot of logarithm ratio of R as a function of log10[Ca 2+ ] (m) at (a)-(f): 10-50°C: [U]total = 50 μm.

Fig. V - 7

 7 Fig. V -7 Plot of logarithm ratio of R as a function of log10[Mg 2+ ] (m) at (a)-(d): 5-30°C: [U]total = 50 μm.

  Fig. V -8 shows that the approaches, including the Van't Hoff equation, the constant heat capacity equation, the DQUANT approach, and the isoelectric modification method, provide satisfactory description of the measured log10K°n.1.3 in the temperature ranges of 5-30°C for MgUO2(CO3)3 2-and 10-50°C for CanUO2(CO3)3 (4-2n)-. However, deviations start to increase at elevated temperature, i.e., 70-90°C, reaching several logarithm units, especially between the values evaluated by the Van't Hoff equation and the constant heat capacity equation. Fig. V -9 presents the variation of log10K°n.1.3 uncertainties bands with temperature using the Van't Hoff approach.

Fig. V - 8

 8 Fig. V -8 Dependence on temperature of the thermodynamic constants for MnUO2(CO3)3 (4-2n)- with lines: the fitted curves based on the Van't Hoff equation, the constant heat capacity equation and the DQUANT equation and red points: the values of log10K°n.1.3 recalculated for each temperature with the isoelectric reaction, as listed in Table V -S7 of the SI.

Fig. V - 9

 9 Fig. V -9 Dependance on temperature of the thermodynamic constants for Mgn/CanUO2(CO3)3 (4-2n)-with lines represented fitted curves based on the Van't Hoff equation, grey short-dashed lines display the 95% confidence interval and red long-dashed lines show the 98% confidence level (see thermodynamic parameters in Table V -2).

  2-in Ca2UO2(CO3)3(aq). This conversion of denticity during the second complexation of CaUO2(CO3)32-by Ca 2+ could partially contribute to an exothermic process due to the bond formation of U-OCO 3 , and a less important entropy value due to the formation of more ordered Ca2UO2(CO3)3(aq), which agrees with our experimental results.The reaction enthalpy and entropy values were determined for the formation of MgUO2(CO3)3 2-from the variation of stability constants in temperature range of 5-30°C with the four approximation approaches. The fitting results of the Van't Hoff equation and the DQUANT equation are of equal magnitude with slightly negative ∆r VH H°m = -(11.653 ± 10.300) kJ mol -1 and ∆r DQUANT H°m = -(11.171 ± 10.683) kJ mol -1 , while the constant heat capacity equation and isoelectric modification result in more exothermic process with ∆r CHC H°m = -(13.477 ± 30.494) kJ mol -1 and ∆r Iso H°m = -(12.025 ± 435) kJ mol -1 . In view of the high uncertainties, the enthalpy contribution to the formation of MgUO2(CO3)3 2-and CaUO2(CO3)3 2- may be both assumed to be almost nil. Our results are approximately consistent with the value directly measured at 25°C as ∆rH°m = (4.3 ± 1.5) kJ mol -1 using calorimetry by Jo, et al.29 It can be concluded that the entropy term contributes primarily to this reaction. However, unlike the comparable values of entropy term obtained for the formations of MgUO2(CO3)3 2-and CaUO2(CO3)3 2-in the work of Jo et al.,28,29 the value of T∆rS°m determined in this study for the complexation of Mg 2+ is globally lower than that of Ca 2+ by about 10-20 kJ mol -1 . The decrease in entropy value from CaUO2(CO3)3 2-to MgUO2(CO3)3 2-might very well be attributed to the more ordered bulk water structure of MgUO2(CO3)3 2-, which would also explain the lower log10K°1.1.3 value for the complexation of Mg 2+ compared to Ca 2+ . Because of its higher charge density, Mg 2+ (charge denticity of 205 C mm -3 calculated from ne/(4/3)πr3 with e = 1.60 10 -19

Fig. V - 10

 10 Fig. V -10 shows the uranium solubility (Fig. V -10 (a)) and speciation (Fig. V -10 (b)) as a function of temperature for the Mol reference water. One can observe that the aqueous system undergoes a fast alteration of dominant species in the narrow temperature range of 40-50°C

  Fig. V -S7 of the SI shows the solubility variation of UO2(cr) as a function of temperature for Mol and COx waters. Compared with Fig. V -10 (a) and Fig. V -11 (a), one can observe that the uranium concentration in solution is lowered by about 10 orders of magnitude in both cases. In equilibrium with UO2(cr), the reduction of U(VI) to U(IV) is more intensely affected by temperature as the evident stronger decrease of the dissolved uranium concentration.

  triscarbonatouranyl(VI) with Mg 2+ /Ca 2+ in 5-30°C for MgUO2(CO3)3 2-and 10-50°C for CanUO2(CO3)3(4-2n)-measured by TRLFS. The variation of log10K°n.1.3 with temperature were fitted by several theoretical approaches: the Van't Hoff equation, the constant heat capacity model, the DQUANT equation, and the isoelectric reaction modification. In the experimental temperature ranges, these approximation models provided satisfactory matching between the experimentally derived constants and the predicted values. Nonetheless, large discrepancy has been observed between the predictions of log10K°n.1.3 by the revised approaches, probably because of the narrow experimental temperature ranges that were limited by the quenching effects with temperature. The Van't Hoff equation and the DQUANT equation showed comparable trends of log10K°n.1.3, while the assumption of non-zero heat capacity resulted in significant deviation at high temperatures. The modification to isoelectric reaction exhibited mildly increasing log10K°n.1.3 with temperature. The accuracy of the experimental values of equilibrium constants at elevated temperatures for these species should be improved by further investigations. The recent thermodynamic data with response to temperature and ionic strength for MnUO2(CO3)3(4-2n)-complexes are highly valuable in estimating the temperature and salinity effects on the speciation of U(VI), especially encountered in nuclear waste repositories and bridging the gaps in the thermodynamic database for the proposed important MnUO2(CO3)3 (4-2n)- complexes.

Fig

  Fig. V -S1. Calibration lines in the temperature range of 5-50°C and the pH values corresponding to the zero point of electrode at 5°C and 50°C.

Fig

  Fig. V -S2 Gaussian decomposition of the luminescence spectrum of the pure spectra of (a) CaUO2(CO3)3 2-and (b) Ca2UO2(CO3)3(aq) normalized to the same area by the MCR-ALS analysis

Fig

  Fig. V -S4 Change in main oxidation states of Fe, S, C, U (a) and variation of pH and pe (b) as a function of temperature for Boom clay equilibrium waters, 11 P(O2) = 10 -68.67 atm using thermodynamic data from the Thermochimie 10a database,12 implementing SrUO2(CO3)3 2-,13 specific ion interaction coefficients of Mgn/CanUO2(CO3)3 (4-2n)-,[4][5][6] and their ΔrH°m values determined in this work based on the isoelectric reaction approximation.

Fig. V -

 V Fig. V -S5 Change in main oxidation states of Fe, S, C, U (a) and variation of pH and pe (b)as a function of temperature for Callovo-Oxfordian clay equilibrium water, 14 P(O2) = 10 -65.83 atm using thermodynamic data from the Thermochimie 10a database,12 implementing SrUO2(CO3)3 2-,13 specific ion interaction coefficients of Mgn/CanUO2(CO3)3 (4-2n)-,[4][5][6] and their ΔrH°m values determined in this work based on the isoelectric reaction approximation.

Fig

  Fig. V -S6 Solubility of UO2:2H2O(am) (a) and aqueous uranium speciation (b) change in main oxidation states of Fe, S, C, U (c) and variation of pH, pe (d), as a function of temperature forCallovo-Oxfordian clay equilibrium waters, 14 P(O2) = 10 -65.83 atm using thermodynamic data from the Thermochimie 10a database,12 implementing SrUO2(CO3)3 2-,13 specific ion interaction coefficients of Mgn/CanUO2(CO3)3(4-2n)-4-6 and their ΔrH°m values determined in this work based on the isoelectric reaction approximation. The reduction reactions of S(+VI) to lower oxidations states were prevented in the input PhreeqC file.

Fig

  Fig. V -S7 Solubility of UO2(cr) as a function of temperature (a) for Boom clay equilibrium waters, 11 P(O2) = 10 -68.67 atm and (b) for Callovo-Oxfordian clay equilibrium waters, 14 P(O2) = 10 -65.83 atm using thermodynamic data from the Thermochimie 10a database,12 implementing SrUO2(CO3)3 2-,13 specific ion interaction coefficients of Mgn/CanUO2(CO3)3(4-2n)-4-6 and their ΔrH°m values determined in this work based on the isoelectric reaction approximation.

Fig. VI - 1

 1 Fig. VI -1 Predominance diagrams obtained using Phreeplot 36 of Mg-Ca-U-CO3 and U-OH calculated for Im = 0.1 mol kgw -1 NaCl using the thermodynamic data and SIT coefficients determined in this work. Calculations performed for C(U)tot = 10 -8 mol kgw -1 , P(O2) from 10 -73 to 10 -51 atm.

Fig. VI - 2 Fig. VI - 3

 23 Fig.VI -2shows the repartition of the reported uranium concentration for the selected water compositions in this study. The measured uranium concentration varies widely from 10 -11 to 5 10 -3 mol kgw -1 . The ionic strength ranges from 0.6 to 83.1 mmol kgw -1 with an average of 12.5 mmol kgw -1 . Some water samples represent the uranium concentration lower than 15 μg L -1

Fig. VI - 4

 4 Fig. VI -4 log10(a(HCO3 -) a(H + )) vs. log10(a(Ca 2+ )/a(H + ) 2 ) predominance plot at [U(VI)] = 100 μmol kgw -1 , P(O2) = 0.2095 atm and fixed values of log10(a(Na + )/a(H + )) = 5. The selected sample solutions having different major species are distinguished by color.

Fig. VI - 5

 5 Fig. VI -5 Repartition of the major species in the speciation calculations of the selected water compositions in a log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(HCO3 -) a(H + ))-log10(a(Na + )/a(H + )) diagram.The values of log10(a(Na + )/a(H + )) were respectively fixed to 2, 5, and 7 to deduce the log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(HCO3 -) a(H + )) predominance diagrams of Ca-UO2-CO3 system at [U(VI)] = 100 μmol kgw -1 , P(O2) = 0.2095 atm.

Fig. VI - 6

 6 Fig. VI -6Major uranium species obtained in the speciation calculations of the selected water compositions of which (a) 3.5 < log10(a(Na + )/a(H + )) < 4.5, (b) 4.5 < log10(a(Na + )/a(H + )) < 5.5, and (c) 5.5 < log10(a(Na + )/a(H + )) < 6.5.

Fig. VI - 7 Fig. VI - 8

 78 Fig. VI -7 focuses on the repartition of the selected water samples in a log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -) a(H + )) diagram with the z-axis varying from -12 to -8.Most of the data points are in the range of -11 < log10(a(HCO3 -) a(H + )) < -9.

Fig. VI - 10

 10 Fig. VI -10 Position of the two Lodève water samples in a log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -) a(H + )) diagram. Ellipsoidal predominance regions of UO2CO3(aq) (blue), CaUO2(CO3)3 2-(black), and Ca2UO2(CO3)3(aq) (red) space determined from the sample point-clouds with the interval confidence of 68%.

  Fig. VI -11 Comparison of luminescence spectra of uranium in a known solution (blue), in the "untreated" water sample (orange) and in the adjusted water at pH 6.5 (grey) from the Lodève site. (a) initial delay time of D = 25 ns and gate width W = 1 μs (b) initial delay time of D =100 ns and gate width W = 1 μs.

Fig

  Fig. VI -14 Comparison of (a) luminescence spectra of uranium in the "untreated" water sample from the Lodève site at 25°C with delay time of D = 15 ns and gate width W = 50 ns before and after long period (≈5h) laser excitation of 450 nm. (b) Normalized spectra to the same area. A set of luminescence decays was recorded for the longtime-excited Lodève water sample at 25°C. As shown in the fitting results in the Fig. VI -15, the luminescence decays biexponentially within the delay range studied. The equation: Intensity = A1exp(-x/tau1) +A2exp(-x/tau2) + y0 was applied in the fitting procedure where A1, A2 represent the luminescence intensity of components 1 and 2 at D = 0, tau1, tau2 (also ꚍ1, ꚍ2) correspond to their decay times and y0 is an adjustable noise term.

Fig

  Fig. VI -15 Bi-exponential fits of luminescence intensity decay from D = 15 ns to 210 ns for Lodève water (a) before and (b) after long period laser excitation of 450 nm

  2+

Table I -

 I

1 Time dimensions for different de-excitation processes.

Table I -

 I 2 Operational conditions and acquisition parameters.

	Optical quartz cuvette	111-QS, 117-QS with screw cap (Suprasil,
		Hellma Analytics)
	Laser source	Laser Nd : YAG (Surelite, Continuum), 10Hz,
		impulse duration : 7 ns
	OPO	Horizon, Continuum
	Spectrometer	Acton 2358i
	CCD camera	Princeton, Andor
	Thermostat	Peter Huber Kältemaschinenbau AG
	Joulemeter	RJP-734, Laser Probe Inc.
	Excitation wavelength	450 nm
	Delay	25 ns
	Gate width	1 μs
	Grating	300, 1800 lines mm -1
	Gain	245
	Accumulation number	1000 for acquisitions at 300 lines mm -1 , 10000 for
		acquisitions at 1800 lines mm -1

Table I

 I

	individual concentration of the three ligands does not vary independently due to the intimately
	connected hydrolysis and carbonate equilibria. For this reason, the uranyl carbonate system
	should be formally understood as a three-component system of uranyl, hydroxide, and
	carbonate.
	Table I -4 Formation constants of uranyl carbonate complexes at standard state conditions
	(Im= 0, T= 298 K) from Guillaumont, et al. 51
	Reaction	log10β °
	UO2 2+ + CO3 2-⇄ UO2(CO3)(aq) 9.94 ± 0.03
	UO2 2+ + 2CO3 2-⇄ UO2(CO3)2 2-16.61 ± 0.09
	UO2 2+ + 3CO3 2-⇄ UO2(CO3)3 4-21.84 ± 0.04
	-4 gives the formation constants of uranyl carbonate complexes used in our study to
	calculate the distribution of aqueous uranyl ion species. Another complication of the aqueous
	carbonate systems is that the hydrolysis reactions are also present, and three ligands -OH -,
	HCO3 -and CO3 2--are always present in carbonated solutions. One cannot investigate metal-
	carbonate reactions without considering the metal ion hydrolysis chemistry because the

Table I -

 I

	0	0.9974	1.0000	\	\	\
	0.1	0.9995	1.0020	0.0649	0.9351	0.9966
	0.2	1.0014	1.0010	0.0725	0.9275	0.9933
	0.5	1.0071	1.0097	0.0723	0.9277	0.9831
	0.75	1.0117	1.0143	0.0656	0.9344	0.9742
	1.0	1.0161	1.0187	0.0568	0.9432	0.9651
	Im (mol L -1 ) NaClO4 ρ (g/cm 3 ) ρ/ρe	1-ϕ	ϕ	aH2O
	0	0.9974	1.0000	\	\	\
	0.1	1.0052	1.0077	0.0672	0.9328	0.9966
	0.5	1.0358	1.0385	0.0841	0.9159	0.9832
	1.0	1.0738	1.0766	0.0808	0.9192	0.9658
	1.5	1.1116	1.1145	0.0735	0.9265	0.9475
	2.0	1.1492	1.1522	0.0625	0.9350	0.9282
	2.2	1.1642	1.1672	0.0613	0.9387	0.9202
	3.0	1.2239	1.2271	0.046	0.9540	0.8866
		ρ (g/cm 3 ) ρ/ρe	1-ϕ	ϕ	aH2O

6

Values of water activity and the osmotic coefficients 116 at investigated ionic strengths in NaCl and NaClO4 used in the calculation of Ringböm coefficients.

Im (mol L -1 ) NaCl

Table I -

 I

	Temp	pH	E(mV)	pH	E(mV)	pH	E(mV)	pH	E(mV)	Slope	Intercept
	5°	1.67	299.1	4.00	166.9	6.95 1.8	9.39	-133.4	-56.01	391.81
	10°	1.67	299.5	4.00	166.2	6.92 0.8	9.33	-135.1	-56.72	393.69
	15°	1.67	302.2	4.00	165.5	6.90 -0.5	9.28	-135.2	-57.45	396.79
	22°	1.68	304.2	4.00	164.9	6.88 -1.4	9.23	-136.1	-58.24	400.18
	30°	1.68	308.2	4.01	165.6	6.85 -3.3	9.14	-138.8	-59.86	407.39
	35°	1.69	311.2	4.02	167.1	6.84 -3.2	9.11	-139.5	-60.70	412.58
	40°	1.69	314	4.03	168.7	6.84 -3	9.07	-140.2	-61.49	417.38
	45°	1.70	318.2	4.04	171.4	6.83 -3.1	9.04	-140	-62.44	423.97
	50°	1.71	321.2	4.06	171.9	6.83 -3.3	9.02	-140.4	-63.16	428.74

7 Recalculated pH values of buffer solutions at investigated temperatures and potential values Emes read on pH-meter.

Table I -

 I 8 Composition and properties of standard buffer solutions.125

	pH value at 25°	1.68	4.01	6.87	9.18
	Composition 125	0.05 mol kgw -1	0.05 mol kgw -1
		potassium	potassium
		tetraoxalate	hydrogen
		(KH3C4O8)	phthalate
			(KHC8H4O4)

  1 mol kgw -1 and of Im = 0.01 mol kgw -1 ; ΔEmes then used in the four-point calibration with commercial buffer solution of pH 1.64, 4.01, 6.87, 9.18, assumed as diluted solutions of Im ≈ 0.01 mol kgw -1 ; the uncertainties of ± 0.05 pH in text included this non-significant experimental bias.

	Solution	Im (mol kgw -1 )	Debye-Hückel term	log10γ	Emes	∆Ej = Emes1 -Emes2	∆Ej -rA0log10(a2/a1)
	HCl 0.01M	0.42
	+ NaCl		0.1674	-0.1662	276.7	22.8	26.0
	0.41 M
	HCl 0.01M	0.30
	+ NaCl		0.1533	-0.1521	276.5	23.0	27.0
	0.29 M
	HCl 0.01M	0.20
	+ NaCl		0.1363	-0.1351	276.4	23.1	28.0
	0.19 M
	HCl 0.01M	0.12
	+ NaCl		0.1150	-0.1137	275.4	24.1	30.2
	0.10 M
	Solution HCl 0.01M	Im (mol kgw -1 ) 0.01	Debye-Hückel term 0.0453	log10γ -0.0441	Emes 275	∆Ej = Emes2 Emes1 -24.5	rA0log10(a2/a1) ∆Ej -34.7
	HCl 0.01M
	+ NaClO4	3.50	0.2498	-0.24812	175.5	0	0
	2.99 M
	HCl 0.01M
	+ NaClO4	2.85	0.2431	-0.2415	168	7.5	7.29
	2.49 M
	HCl 0.01M
	+ NaClO4	2.22	0.2343	-0.2328	163.6	11.9	11.56
	1.99 M
	HCl 0.01M
	+ NaClO4	1.62	0.2225	-0.2210	158.9	16.6	16.269
	1.49 M
	HCl 0.01M
	+ NaClO4	1.05	0.2057	-0.2042	155.1	20.4	20.39
	0.99 M
	HCl 0.01M
	+ NaClO4	0.51	0.1750	-0.1735	150.2	25.3	26.32
	0.49 M
	HCl 0.01M
	+ NaClO4	0.20	0.1392	-0.1378	146.6	28.9	31.81
	0.19 M
	HCl 0.01M	0.01	0.0441	-0.0427	130.2	45.3	53.32
	HCl 0.01M	3.19
	+ NaCl		0.2471	-0.2458	299.5	0.0	0
	2.97 M
	HCl 0.01M	1.56
	+ NaCl		0.2212	-0.2200	286	13.5	14.2
	1.49 M
	HCl 0.01M	1.04
	+ NaCl		0.2051	-0.2039	280.8	18.7	20.0
	0.99 M
	HCl 0.01M	0.76
	+ NaCl		0.1924	-0.1911	279.7	19.8	21.7
	0.74 M
	HCl 0.01M	0.61
	+ NaCl		0.1832	-0.1820	277.9	21.6	24.0
	0.59 M
	HCl 0.01M	0.51
	+ NaCl		0.1754	-0.1742	277.8	21.7	24.5
	0.49 M

  4-contributed primarily to the variation of non-complexed by calcium species as a function of pH, over other hydroxides and carbonates species in calcium-uranium(VI)carbonate system. Consequently, evaluation of the Ringböm44 coefficient α at different pH values in free-calcium conditions sufficed to estimate the impact of pH values. The assumption is then validated by speciation calculation with determined thermodynamic data in this work.

Table II -

 II

	Reaction	log10 K °	References
	CO2(g) + H2O ⇄ CO3 2-+ 2H +	-18.15 ± 0.06	32
	UO2 2+ + H2O ⇄ UO2(OH) + + H +	-5.25 ± 0.24	32
	UO2 2+ + 2H2O ⇄ UO2(OH)2(aq) + 2H +	-12.15 ± 0.07	32
	3 UO2 2+ + 5H2O ⇄ (UO2)3(OH)5 + + 5H +	-15.55 ± 0.12	32
	UO2 2+ + CO3 2-⇄ UO2(CO3)(aq)	9.94 ± 0.03	32
	UO2 2+ + 2 CO3 2-⇄ UO2(CO3)2 2-	16.61 ± 0.09	32
	UO2 2+ + 3 CO3 2-⇄ UO2(CO3)3 4-	21.84 ± 0.04	32
	2 UO2 2+ + CO2(g) + 4H2O ⇄ (UO2)2CO3(OH)3 -+ 5H +	-19.01 ± 0.50	32
	UO2(CO3)3 4-+ 4H + ⇄ UO2(CO3)(aq) + 2CO2(g) + 2H2O	24.40 ± 0.21*	Recalculated from 32
	UO2(CO3)3 4-+ 2H + ⇄ UO2(CO3)2 2-+ CO2(g) + H2O	12.92 ± 0.15*	Recalculated from 32
	2 UO2(CO3)3 4-+ 7 H + ⇄ (UO2)2CO3(OH)3 -+ 5CO2(g) + 2H2O	46.21 ± 0.70*	Recalculated from 32
	Ca 2+ + UO2(CO3)3 4-⇄ CaUO2(CO3)3 2-	5.26 ± 0.01	This work
	2Ca 2+ + UO2(CO3)3 4-⇄ Ca2UO2(CO3)3(aq)	8.65 ± 0.02	This work
	Ca 2+ + UO2 2+ + 3 CO3 2-⇄ CaUO2(CO3)3 2-	27.20 ± 0.04	This work
	2Ca 2+ + UO2 2+ + 3 CO3 2-⇄ Ca2UO2(CO3)3(aq)	30.49 ± 0.05	This work

1 Stability constants of U(VI) complexes at 25 °C and Im = 0, used in the work.

Table II -

 II 2 Gibbs energies of formation from the main species considered within this work.

	Species	ΔfG°m (kJ mol -1 )	References
	H2O	-237.140 ± 0.041	32
	CO2(g)	-394.373 ± 0.133	32
	CO3 2-	-527.900 ± 0.390	32
	UO2 2+	-952.551 ±1.747	32
	UO2(CO3)3 4-	-2 660.914 ± 2.116	32
	(UO2)2CO3(OH)3 -	-3 139.526 ± 4.517	32
	Ca 2+	-552.806 ± 1.050	32
	CaUO2(CO3)3 2-	-3 244.316 ± 2.361	This work
	Ca2UO2(CO3)3(aq)	-3 815.901 ± 2.985	This work

Table II -

 II 3 Main specific ion interaction coefficients used in this work.

	Specific ion interaction coefficient	Value (± 1σ)	Reference
	ε(H + ,Cl -)	0.12 ± 0.01	32
	ε(Na + ,CO3 2-)	-0.08 ± 0.03	32
	ε(UO2 2+ ,Cl -)	0.21 ± 0.02	32
	ε(Ca 2+ ,Cl -)	0.14 ± 0.01	32
	ε(UO2(CO3)2 2-,Na + )	-0.02 ± 0.09	32
	ε(UO2(CO3)3 4-,Na + )	-0.01 ± 0.11	32
	ε(CaUO2(CO3)3 2-,Na + )	0.29 ± 0.11	This work
	ε(Ca2UO2(CO3)3(aq),NaCl)	0.66 ± 0.12	This work

  -4 Values of decimal logarithm of calcium concentration (mol kgw -1 ) extracted from the slope analysis in Fig.II -4, and pH values measured before luminescence acquisition.

	Im (mol kgw -1 )	CaUO2(CO3)3 2-	Ca2UO2(CO3)3(aq)
	0.10	-3.35 ≤ log10[Ca 2+ ] ≤ -2.58	-2.58 ≤ log10[Ca 2+ ] ≤ -1.52
		8.10 ≤ pH ≤7.85	7.85 ≤ pH ≤7.45
	0.20	-3.00 ≤ log10[Ca 2+ ] ≤ -2.15	-2.15 ≤ log10[Ca 2+ ] ≤ -1.41
		8.14 ≤ pH ≤7.75	7.75 ≤ pH ≤7.40
	0.50	-2.50 ≤ log10[Ca 2+ ] ≤ -1.87	-1.87 ≤ log10[Ca 2+ ] ≤ -1.43
		7.91 ≤ pH ≤7.68	7.68 ≤ pH ≤7.49
	0.75	-2.20 ≤ log10[Ca 2+ ] ≤ -1.44	-1.44 ≤ log10[Ca 2+ ] ≤ -1.07
		7.79 ≤ pH ≤7.52	7.52 ≤ pH ≤7.39
	1.00	-2.12 ≤ log10[Ca 2+ ] ≤ -1.40	-1.40 ≤ log10[Ca 2+ ] ≤ -1.08
		7.86 ≤ pH ≤7.54	7.54 ≤ pH ≤7.40
	The binding constants of Ca 2+ ion to CaUO2(CO3)3 2formation constants by extrapolating to infinite dilution using the SIT to log10K°1.1.3 =
	5.36 ± 0.02 and log10K°2.1.3 = 8.65 ± 0.02. For CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), the
	cumulative formation constants are log10°1.1.3 = 27.20 ± 0.04 and log10°2.1.3 = 30.49 ± 0.05,

-, noted as ∆log10K, are 2.43, 2.13, 1.79, 1.59, and 1.51 by subtracting log10K1.1.3 from log10K2.1.3. It is found that ∆log10K is lower than log10K1.1.3 at each ionic strength, revealing that the binding of Ca 2+ to CaUO2(CO3)3 2-is more difficult than that of Ca 2+ to UO2(CO3)3 4-. This is expected because the electrostatic interaction is stronger between Ca 2+ /UO2(CO3)3

4-

than that between Ca 2+ /CaUO2(CO3)3 2-due to the stepwise charge neutralization.

The resulting conditional equilibrium constants after rounding off the slope values to their nearest integers are given in Table

II

-S3 of the SI. By comparing the different ionic strength series, one can immediately assess the tendency of the stability of ternary species. The dependence of log10K1.1.3 + 16D and log10K2.1.3 + 24D vs. the molal ionic strength in NaCl is represented in Fig.

II -5

. The two plots both give a linear relationships over the entire range of ionic strength, which allow determining the ion interaction coefficients and the stepwise respectivelysee Table

II

-1, and see ΔfG°m in Table

Table II -

 II S1 Aqueous conditions in preliminary calculations.

	Sample No.	pH	log10([

Ca 2+ ] /mol kgw -1 ) [Ca 2+ ] (mol kgw -1 )

	8 Sample No.	pH	log10([Ca 2+ ] /mol kgw -1 )	[Ca 2+ ] (mol kgw -1 )	Dolomite
	21		7.65	-2.10		7.94 х 10 -3	Calcite
	22 23	6	7.60 7.55	-2.00 -1.90		1.00 х 10 -2 1.26 х 10 -2	Nesquehonite
	24		7.50	-1.80		1.58 х 10 -2
	25 26	4	7.45 7.40	-1.70 -1.60	Dolomite	2.00 х 10 -2 2.51 х 10 -2
	27		7.35	-1.50		3.16 х 10 -2
	28 29	2	7.30 7.25	-1.40 -1.30		3.98 х 10 -2 5.01 х 10 -2
	30		7.20	-1.20		6.31 х 10 -2
	31 SI 32	0	7.15 7.10	-1.10 -1.00		7.94 х 10 -2 1.00 х 10 -1
	33 34	-2	7.05 Calcite 7.00	-0.90 -0.80		1.26 х 10 -1 1.58 х 10 -1
		-4		Nesquehonite
		-6
		-8
		6	7	8		9
				pH
	1		9.00	---		1.00 × 10 -5
	2		8.90	-4.90		1.26 × 10 -5
	3		8.80	-4.70		2.00 × 10 -5
	4		8.70	-4.50		3.16 × 10 -5
	5		8.60	-4.30		5.01 × 10 -5
	6		8.50	-4.10		7.94 × 10 -5
	7		8.40	-3.90		1.26 × 10 -4
	8		8.30	-3.70		2.00 × 10 -4
	9		8.25	-3.50		3.16 × 10 -4
			8.20	-3.30		5.01 × 10 -4
			8.15	-3.10		7.94 × 10 -4
			8.10	-3.00		1.00 × 10 -3
			8.05	-2.90		1.26 × 10 -3
			8.00	-2.80		1.58 × 10 -3
			7.95	-2.70		2.00 × 10 -3
			7.90	-2.60		2.51 × 10 -3
			7.85	-2.50		3.16 × 10 -3
			7.80	-2.40		3.98 × 10 -3
			7.75	-2.30		5.01 × 10 -3
			7.70	-2.20		6.31 × 10 -3

Table II -

 II S2 Calculated Ringböm coefficients α for experimental samples.

		Im = 0.1 mol kgw -	Im = 0.2		Im = 0.5		Im = 0.75		Im = 1 mol kgw -1
		1		mol kgw -1		mol kgw -1		mol kgw -1
	Sample	pH	α	pH	α	pH	α	pH	α	pH	α
		value		value		value		value		value
		9.00	1.00	8.98	1.00	9.00	1.00	9.05	1.00	8.94	1.00
		8.66	1.06	8.45	1.00	8.42	1.01	8.45	1.00	8.40	1.01
		8.42	1.04	8.28	1.02	8.32	1.03	8.24	1.01	8.44	1.04
		8.26	1.10	8.21	1.04	8.17	1.04	8.10	1.01	8.12	1.03
		8.23	1.12	8.18	1.04	8.12	1.02	8.02	1.02	8.04	1.03
		8.15	1.24	8.16	1.06	8.10	1.03	7.90	1.06	7.85	1.04
		8.11	1.37	8.14	1.06	8.07	1.05	7.83	1.12	7.77	1.13
		8.10	1.34	8.13	1.07	8.02	1.05	7.79	1.19	7.83	1.06
		8.08	1.51	8.05	1.10	7.91	1.11	7.75	1.29	7.73	1.20
		8.00	2.06	7.93	1.59	7.89	1.15	7.70	1.53	7.66	1.46
		7.97	2.52	7.90	1.88	7.83	1.27	7.62	2.28	7.60	1.85
		7.92	3.30	7.85	2.42	7.76	1.63	7.57	2.94	7.55	2.38
		7.87	4.79	7.82	2.90	7.72	1.97	7.53	3.94	7.52	3.04
		7.85	5.51	7.77	4.05	7.69	2.31	7.50	4.85	7.48	3.86
		7.81	7.07	7.75	4.73	7.67	2.57	7.45	6.95	7.36	9.08
		7.75	11.55	7.73	5.42	7.66	2.73	7.40	10.15	7.32	12.31
		7.63	28.25	7.64	10.60	7.65	2.91	7.35	14.87	7.28	16.98
		7.60	37.08	7.57	18.27	7.60	4.01	7.31	20.34	--	--
		7.55	53.34	7.54	23.09	7.54	6.16	7.28	25.78	--	--
		7.52	70.02	7.44	50.93	7.49	8.89	7.24	35.42	--	--
		7.50	79.91	7.36	98.01	7.41	16.67	--	--	--	--
		7.48	96.30	--	--	7.38	21.02	--	--	--	--

Table II -

 II S3 Stepwise formation constants derived from the rounded off slopes through the linear dependence of log10R on log10[𝐶𝑎 2+ ] (mol•kgw-1 ).

	Chapter III. Spectroluminescence	Measurements	of	Stability
	Constants of Ca n UO 2 (CO 3 ) 3 (4-2n)-Complexes in NaClO 4 Medium and
	Investigation of Interaction Effects.
	Chengming Shang, Pascal E. Reiller* and Thomas Vercouter
	1 Den -Service d'Études Analytiques et de Réactivité des Surfaces (SEARS), CEA, Université Paris-Saclay,
	F-91191, Gif-sur-Yvette, France
	*Corresponding author. Tel.: +33 1 6908 4312; fax: +33 1 6908 9475. E-mail address: pascal.reiller@cea.fr
	Dalton Transactions 49 (43), 15443-15460.
		http://doi.org/10.1039/D0DT03164J
	Im (mol•kgw -1 )	log10K113		log10K213
	0.10	3.58 ± 0.02	6.01 ± 0.03
	0.20	3.17 ± 0.02	5.30 ± 0.03
	0.50	2.50 ± 0.01	4.29 ± 0.02
	0.75	2.16 ± 0.02	3.75 ± 0.01
	1.00	1.95 ± 0.02	3.41 ± 0.03

2

  𝑙𝑜𝑔 10 𝛽 𝑛.1.3 = 𝑙𝑜𝑔 10 𝛽 0.1.3 + 𝑙𝑜𝑔 10 𝐾 𝑛.1.3 = 𝑙𝑜𝑔 10 𝛽 𝑛.1.3 ° + 𝛥𝑧 2 𝐷 𝐻 -𝛥𝜀 𝐼 𝑚 (III -6)

  NaClO4 with between 14 and 19 samples per series (Table III -S1of the SI). Out of 114 luminescence spectra, 96 were used to estimate the specific ion interaction

	1, 0.51, 1.05,
	1.62, 2.21, 2.46, and 3.5 m

Table III

 III

	-3 Values of log10[Ca 2+ ] extracted from the slope analysis in Fig. III -2, and pH
	values measured before luminescence acquisition.
	Im	CaUO2(CO3)3 2-(log10 [Ca 2+ ], pH)	Ca2UO2(CO3)3(aq) (log10 [Ca 2+ ], pH)
	0.10 m	(-3.60, 8.21) → (-2.71, 7.87)	(-2.51, 7.80) → (-1.87, 7.46)
	0.51 m	(-2.80, 7.84) → (-2.12, 7.64)	(-2.00, 7.60) → (-1.66, 7.44)
	1.05 m	(-2.67, 7.78) → (-2.11, 7.56)	(-1.94, 7.52) → (-1.60, 7.34)
	1.62 m	(-2.38, 7.72) → (-1.88, 7.52)	(-1.79, 7.48) → (-1.43, 7.32)
	2.21 m	(-2.15, 7.82) → (-1.78, 7.50)	(-1.68, 7.48) → (-1.26, 7.22)
	2.46 m

(-2.21, 7.58

) → (-1.70, 7.34) (-1.58, 7.30) → (-1.

30, 7.18)

Table III -

 III

	Species	ΔfG°m kJ mol -1	Ref.
	H2O	-237.140 ± 0.041	29
	CO2(g)	-394.373 ± 0.133	29
	CO3 2-	-527.900 ± 0.390	29
	UO2 2+	-952.551 ± 1.747	29
	UO2(CO3)3 4-	-2 660.914 ± 2.116	29
	Ca 2+	-552.806 ± 1.050	29
	CaUO2(CO3)3 2-	-3 244.316 ± 2.361	8
		-3 244.658 ± 2.394	p.w.
		-3 244.430 ± 2.361	p.w. multilinear regression
	Ca2UO2(CO3)3(aq)	-3 815.901 ± 2.985

5

Gibbs energies of formation from the main species considered.

  -2. The sets of data were satisfactory fitted within the SIT framework. The stability constants at infinite dilution were virtually the same in both media, whatever the hypothesis considering the uncertainties. The increasing tendency of log10 2.1.3 + 24DH vs. Im in NaClO4 is comparable with that reported by Kalmykov and Choppin.2 Conversely, the slopes that represents -∆ε, are different in NaCl and

Table III -

 III 6 Aqueous composition in binary/ternary species for specifically selected sample solutions, of which luminescence emission spectra are shown in Fig.III -5.

			UO2(CO3)3 4-	CaUO2(CO3)3 2-	Ca2UO2(CO3)3(aq)
	Im (m)	(a)	(b)	(a)	(b)	(a)	(b)
	0.10	2%	2%	43%	37%	50%	55%
	0.51	7%	12%	42%	50%	42%	29%
	1.05	6%	15%	31%	51%	42%	30%
	1.62	8%	25%	38%	45%	48%	26%
	2.21	8%	18%	35%	66%	44%	16%
	2.46	9%	23%	36%	66%	41%	10%

Table III -

 III

	Ref.	SIT Eq. III -4					Ciavatta 73 Eq. III -11
		log10 β°n.1.3	Δε(NaClO4)	RSS* log10 β°n.1.3	ε1	ε2	RSS*
	CaUO2(CO3)3 2-	27.26 ±	-0.47 ± 0.04	19.97	27.02 ±	-0.81	0.67 ±	1.86
		0.08				0.05	± 0.06	0.12
	Ca2UO2(CO3)3(aq)	30.53 ±	-0.58 ± 0.07	40.10	30.21 ±	-1.03	0.92 ±	10.42
		0.12				0.13	± 0.16	0.32
		SIT multilinear
		log10 β°n.1.3	Δε(NaClO4)
				Δε(NaCl)
	CaUO2(CO3)3 2-	27.22 ±	-0.49 ± 0.03	-0.04 ±
		0.04	0.02
	Ca2UO2(CO3)3(aq)	30.51 ±	-0.59 ± 0.04	0.16 ±
		0.06	0.12
	Ref.		Bretti et al. 84 Eq. III -9		Bretti et al. 85 Eq. III -10
		log10 β°n.1.3	ε	K	RSS* log10 β°n.1.3	ε0	ε∞	RSS*
	CaUO2(CO3)3 2-	27.02 ±	-0.14 ±	3.63 ± 1.18	2.00	27.03 ±	-1.40	-0.19 ±	1.76
		0.07	0.03			0.05	± 0.17	0.05
	Ca2UO2(CO3)3(aq)	30.18 ±	-0.21 ±	5.78 ± 4.71	11.79	30.23 ±	-1.83	-0.19 ±	10.40
		0.21	0.08			0.12	± 0.43	0.14
	* Residual Sum of Square RSS = ∑ w i (y i -y i ̂)2 n i=1	, with wi = 1/σ². The values of σi are the uncertainties associated with
	log10 β1.1.3 + 16DH for CaUO2(CO3)3 2-, log10 β2.1.3 + 24DH for Ca2UO2(CO3)3(aq) in Table III -4 and y i ̂ are estimated values
	by weighted least squares regression for the datasets (Im, log10 β1.1.3 + 16DH /log10 β2.1.3 + 24DH).

7

Parameter values by fitting the experimental results at ionic strengths from 0.1 to 2.46 m NaClO4 to the models under consideration.

Table III -

 III S1 Potential values Emes read on pH-meter for NaClO4 solutions of Im = 0.01 to 3.5 m at pcH = 2 using the reconditioned electrode of 3.5 m NaClO4 reference solution: liquid junction potentials ∆Emes calculated to estimate the potential differences between the solutions of a specific Im ≥ 0.1 m and of Im = 0.01 m; ∆Emes then used in the four-point calibration with commercial buffer solution of pH 1.64, 4.01, 6.87, 9.18, assumed as diluted solutions of Im ≈ 0.01 m; the uncertainties of ± 0.05 pH in text included this non-significant experimental bias.

	2		8.45 8.30	1.02 1.01	5.03E-05 0.00097	1.75E+08 2.39E+08	8.13	1.02	3.44E-04	2.10E+08
	3		8.41 8.10	1.03 1.03	9.51E-05 0.00306	1.55E+08 3.87E+08	7.95	1.09	8.46E-04	1.95E+08
	4		8.33 7.90	1.06 1.03	1.35E-04 0.00719	1.44E+08 1.92E+08	7.90	1.17	1.06E-03	1.95E+08
	5		8.25 7.70	1.12 1.06	1.58E-04 0.01769	2.23E+08 2.28E+08	7.87	1.24	1.24E-03	2.10E+08
	6		8.21 7.60	1.17 1.11	2.49E-04 0.02156	1.48E+08 7.84 1.86E+08	1.27	1.59E-03	2.05E+08
	7		8.17 7.50	1.19 1.29	2.88E-04 0.02623	1.41E+08 7.81 2.04E+08	1.38	2.00E-03	2.23E+08
	8		8.13 7.44	1.29 1.49	4.36E-04 0.03109	1.76E+08 7.78 2.08E+08	1.54	2.39E-03	2.24E+08
	9		8.09 7.39	1.45 1.85	5.06E-04 0.04189	1.61E+08 7.75 2.18E+08	1.73	2.99E-03	2.40E+08
	10		8.01 7.33	2.03 2.58	7.77E-04 0.05419	1.77E+08 7.72 2.27E+08	2.00	3.86E-03	2.36E+08
	11		7.91 7.29	3.69 3.33	1.13E-03 0.06825	2.03E+08 7.69 2.07E+08	2.36	4.45E-03	2.58E+08
	12		7.87 7.27	4.84 3.84	1.94E-03 0.09060	1.92E+08 7.66 2.48E+08	2.83	5.20E-03	2.72E+08
	13		7.80 7.25	8.04 3.08E-03 4.44 0.14429	1.98E+08 7.64 2.79E+08	3.21	7.61E-03	2.79E+08
	14		7.73 7.22	13.67 5.05E-03 5.56 0.21003	2.06E+08 7.60 3.31E+08	4.21	9.97E-03	2.97E+08
	15		7.68 7.20	20.17 6.27E-03 6.51 0.25634	2.27E+08 7.56 3.57E+08	5.59	1.27E-02	3.06E+08
	16		7.62 7.17	32.36 8.69E-03 8.23 0.32939	2.33E+08 7.52 3.88E+08	7.54	1.56E-02	3.41E+08
	17		756 7.16	52.16 1.00E-02 9.00 0.39816	2.30E+08 7.48 4.09E+08	10.22	1.69E-02	3.63E+08
	18		7.48 7.15	98.99 1.26E-02 9.85 0.50482	2.35E+08 7.44 4.05E+08	13.97	2.20E-02	3.60E+08
			7.46	116.8	1.34E-02	2.40E+08
	Im (mol•kgw -1 ) 1.05							1.62
	Sample		pH	α		[Ca 2+ ]			F0	pH	α	[Ca 2+ ]	F0 (counts)
			value		(mol•kgw -	(counts)	value		(mol•kgw -1 )
						1 )
			9.00	1.00		0		2.27E+08	9.00 1.00	0	5.45E+08
	Solution		8.13 7.95 7.80 7.78	Conversion factor (from molarity M to 1.01 8.81E-04 Im (mol kgw -1 ) 1.71E+08 1.04 1.46E-03 1.55E+08 1.08 1.94E-03 1.54E+08 molality m) 1.11 2.16E-03 1.73E+08 7.72 1.22 Emes (mV) 7.95 1.04 7.80 1.09 7.78 1.12	5.53E-04 ΔEmes = Emes -Emes 4.01E+08 1.59E-03 3.91E+08 (Im = 0.01 m) 2.97E-03 3.75E+08 4.15E-03 3.71E+08
	HCl 0.01 M		7.72	1.1678 1.20	3.50 2.70E-03 1.69E+08 7.68 1.36 172.2	31.3 4.61E-03	4.19E+08
	+ NaClO4 2.99 M	7.68	1.32	3.45E-03	1.70E+08 7.64 1.56	5.58E-03	4.00E+08
	HCl 0.01 M		7.64	1.1361 1.51	2.84 4.32E-03 2.33E+08 7.60 1.88 168	27.1 6.65E-03	4.22E+08
	+ NaClO4 2.49 M	7.60	1.80	5.57E-03	2.15E+08 7.56 2.33	8.82E-03	4.42E+08
	HCl 0.01 M + NaClO4 2.19 M HCl 0.01 M + NaClO4 1.99 M HCl 0.01 M	7.56 7.52 7.48 7.45 7.40 7.34	1.1196 2.22 2.81 1.14E-02 2.46 7.77E-03 2.35E+08 7.52 2.97 165.5 2.24E+08 7.48 3.86 1.1062 2.21 163.6 3.64 1.33E-02 2.57E+08 7.45 4.79 4.50 1.64E-02 2.60E+08 7.40 6.95 1.0780 1.62 158.9 6.52 2.18E-02 2.80E+08 7.34 11.08 10.36 2.53E-02 3.53E+08 7.32 15.00	1.31E-02 24.6 1.64E-02 2.25E-02 22.7 2.83E-02 3.45E-02 18 3.73E-02	4.68E+08 5.07E+08 5.05E+08 5.61E+08 6.18E+08 6.13E+08
	+ NaClO4 1.49 M Im (mol•kgw -1 ) 2.21							2.46
	HCl 0.01 M Sample		pH	1.0515 α		[Ca 2+ ]	1.05	F0	pH	155.1 α	[Ca 2+ ]	14.2	F0 (counts)
	+ NaClO4 0.99 M HCl 0.01 M + NaClO4 0.49 M HCl 0.01 M + NaClO4 0.19 M	value 9.06 7.85 7.82 7.79	1.0265 1.00 1.0098 1.02 1.05 1.05	(mol•kgw -0.51 1 ) 0 3.02E+08 (counts) 0.20 3.72E-03 2.81E+08 7.07E-03 3.14E+08 7.95 1.01 value 150.2 8.80 1.00 146.6 8.13 1.00 7.80E-03 3.37E+08 7.80 1.02	(mol•kgw -1 ) 9.3 0 2.45E-03 5.7 3.93E-03 4.99E-03	4.77E+08 4.61E+08 4.34E+08 4.30E+08
	HCl 0.01 M		7.73	1.0075 1.07	0.10 8.77E-03 3.53E+08 7.58 1.17 144.5	3.6 6.20E-03	3.96E+08
	+ NaClO4 0.09 M	7.68	1.11	9.67E-03	3.98E+08 7.54 1.28	7.74E-03	3.91E+08
	HCl 0.01 M		7.62	1.0000 1.19	0.01 1.16E-02 4.40E+08 7.50 1.45 140.9	0 9.51E-03	4.47E+08
	For the solution of Im < 0.1 M, the conversion factor taken as 1. The factor values for Im ≥ 0.1 M taken from 7.58 1.24 1.51E-02 4.42E+08 7.46 1.69 1.13E-02 4.17E+08
	7.50 Guillaumont et al. 1 7.48	1.60 1.75 2.11E-02 1.68E-02	4.28E+08 7.42 2.06 4.16E+08 7.38 2.60	1.27E-02 1.72E-02	4.06E+08 4.65E+08
	7.42 Table III -S2 Experimental samples information -pH value, calculated Ringböm coefficients 2.38 2.89E-02 4.41E+08 7.34 3.34 2.01E-02 5.22E+08 7.35 3.71 3.27E-02 4.59E+08 7.30 4.42 2.63E-02 5.52E+08 α, [Ca 2+ ] (mol•kgw -1 ) and deduced F0. 7.30 5.29 4.24E-02 5.05E+08 7.27 5.50 3.46E-02 5.75E+08
			7.22	9.62 5.54E-02	5.56E+08 7.22 7.97	4.13E-02	6.82E+08
	Im (mol•kgw -1 ) 0.10							0.51 7.18 10.85	5.01E-02	7.59E+08
	Sample Im (mol•kgw -1 ) 3.50 pH	α	[Ca 2+ ]		F0		pH	α	[Ca 2+ ]	F0 (counts)
	Sample	value pH value	α	[Ca 2+ ] (mol•kgw -(mol•kgw -	(counts) F0 (counts)	value		(mol•kgw -1 )
					1 )	1 )
	1		9.00 8.94	1.00 1.00	0 0.00001		1.49E+08 9.00 1.28E+08	1.00	0	2.48E+08

Table III -

 III S3 Information for samples in Fig.III -4.

		0.1 m NaCl	0.1 m NaClO4	1 m NaCl	1 m NaClO4
	pH value	7.82	7.78	7.43	7.56
	[Ca 2+ ] (mol kgw -1 )	5.98E-3	5.19E-3	7.80E-2	2.80E-2

Table III

 III

	Im (mol•kgw -1 )	τ (UO2(CO3)3 4-) (ns)	τ (CaUO2(CO3)3 2-) (ns)	τ (Ca2UO2(CO3)3(aq)) (ns)
	0.10	10.62 ± 0.22	30.41 ± 1.9	51.19 ± 1.85
	0.51	8.46 ± 0.17	39.79 ± 2.78	68.53 ± 4.45
	1.05	9.51 ± 0.11	37.57 ± 4.77	63.02 ± 4.45
	1.62	9.39 ± 0.13	35.52 ± 2.66	65.76 ± 6.35
	2.21	11.21 ± 0.40	27.72 ± 1.34	47.54 ± 2.53
	2.46	10.61 ± 0.12	38.6 ± 4.33	86.36 ± 10.15
	Averaged τ	10.0 ± 1.0 (τ measured	34.20 ± 2.24 (τ measured	59.21 ± 4.15 (τ measured from
		from Im = 0.1 to 2.46 m)	from Im = 0.1 to 2.21 m)	Im = 0.1 to 2.21 m)

-S4 Decay-times of CanUO2(CO3)3

(4-2n)

-determined at each Im.

Table IV -

 IV 3 Main specific ion interaction coefficients used in this work.

	Specific ion interaction coefficient	Value (± 1σ)	Ref.
	ε(H + ,ClO4 -)	0.14 ± 0.02	Guillaumont, et al. 22
	ε(H + ,Cl -)	0.12 ± 0.01	idem
	ε(Na + ,CO3 2-)	-0.08 ± 0.03	idem
	ε(UO2 2+ ,ClO4 -)	0.46 ± 0.03	idem
	ε(UO2 2+ ,Cl -)	0.46 ± 0.03	idem
	ε(Mg 2+ ,ClO4 -)	0.33 ± 0.03	idem
	ε(Mg 2+ ,Cl -)	0.19 ± 0.02	idem
	ε(UO2(CO3)2 2-,Na + )	-0.02 ± 0.09	idem
	ε(UO2(CO3)3 4-,Na + )	-0.01 ± 0.11	idem
	ε(MgUO2(CO3)3 2-,Na + ) (in NaClO4)	0.19 ± 0.11	p.w.
	ε(MgUO2(CO3)3 2-,Na + ) (in NaCl)	0.09 ± 0.16	p.w.
	ε(CaUO2(CO3)3 2-,Na + ) (in NaCl)	0.29 ± 0.11	Shang and Reiller 24
	ε(Ca2UO2(CO3)3,NaCl)	0.66 ± 0.22	idem
	ε(CaUO2(CO3)3 2-,Na + ) (in NaClO4)	0.02 ± 0.04	idem
	ε(Ca2UO2(CO3)3,NaClO4)	0.18 ± 0.07	idem

Table IV -

 IV 5 Formation constants log10 K1.1.3 derived from the rounded off slopes in Fig. IV -1. Values of log10 β0.1.3 were calculated with log10 β0.1.3° = 21.84 and Δε = -0.23. kgw mol -1 .

	Im (NaClO4)	log10 K1.1.3	log10 β0.1.3	log10 β1.1.3 + 16DH	Im (NaCl)	log10 K1.1.3	log10 β0.1.3	log10 β1.1.3 + 16DH
	0.10 m	2.61 ±	21.86 ±	26.23 ± 0.05	0.10 m	2.98 ±	21.86 ±	26.59 ± 0.05
	0.51 m	0.03 1.69 ±	0.04 21.96 ±	26.46 ± 0.04	0.25 m	0.02 2.35 ±	0.04 21.90 ±	26.58 ± 0.04
	1.05 m	0.02 1.40 ±	0.04 22.08 ±	26.77 ± 0.05	0.50 m	0.02 1.88 ±	0.04 21.96 ±	26.63 ± 0.05
	1.62 m	0.02 1.26 ±	0.04 22.21 ±	27.04 ± 0.04	0.75 m	0.03 1.56 ±	0.04 22.01 ±	26.64 ± 0.04
	2.21 m	0.02 1.10 ±	0.04 22.35 ±	27.20 ± 0.04	1.00 m	0.01 1.33 ±	0.04 22.07 ±	26.66 ± 0.04
		0.02	0.04			0.02	0.04

Table IV -

 IV 6 Formation constants of MgUO2(CO3)3 2-complex determined in literature and in this work.

	log10β°(MgUO2(CO3)3 2-)	Medium and Im correction	Ref
	26.11 ± 0.04	0.1 mol L -1 NaNO3 (Davies)	Dong and Brooks 14
	25.8 ± 0.5	Averaged from values in 0.101 to 0.509 mol kgw -1 NaNO3 (Davies) Dong and Brooks 15
	25.02 ± 0.08	Extrapolated using SIT 0.101 -0.509 mol kgw -1 NaNO3	Dong and Brooks 15
	26.25 ± 0.04	0.1 mol L -1 NaCl (SIT)	Endrizzi and Rao 18
	25.8 ± 0.3	0.15 mol L -1 NaCl (SIT)	Lee, et al. 10
	26.24 ± 0.13	0.1 mol L -1 NaClO4 (Davies)	Geipel, et al. 12
	26.57 ± 0.01	0.1 -1.0 mol kgw -1 NaCl (SIT)	p.w.
	26.27 ± 0.06	0.1 -2.21 mol kgw -1 NaClO4 (SIT)	p.w.

  Table IV -S3 of SI for detailed decomposition parameters). The variations of characteristic peak wavelengths are weak and stay within the uncertainties when favouring the formations of ternary species. The spectral peak positions are then weightily averaged with 2σ uncertainties and found at 464.98 ± 0.44

	nm, 484.77 ± 0.22 nm, 503.50 ± 0.49 nm, 524.84 ± 1.11 nm and 542.28 ± 2.71 nm in NaClO4
	and 464.63 ± 0.32 nm, 483.77 ± 0.29 nm, 503.71 ± 0.33 nm, 525.08 ± 0.80 nm and 546.24 ±
	2.21 nm in NaCl. No obvious shifts in characteristic peak positions emerged in Mg-UO2-CO3

Table IV -

 IV S10 of the SI) and suggest that the use of log10K°n.1.3 or log10β°n.1.3 and the

	formalism of activity corrections did not generate too much of an inconsistency in this case
	when the calculations have the same source of formation constants for
	MnUO2(CO3)3 (4-2n)-species.

  The conversion factor was taken as 1 for the solution of Im = 0.01 M,. The factor values for the solutions of Im ≥ 0.1 M were either a taken from Guillaumont et al.1 or b extrapolated from these known values.

	Solution	Conversion	factor	Im (mol kgw -1 ) Emes	∆Emes = Emes (Im = 3.19 m)
		(from molarity M to		(mV)	-Emes
	Number of tables: 10	molality m)
	Number of figures: 6 HCl 0.01M + NaCl 0.19M	1.0040 b		0.20	276.4	23.1
	Table IV -S1 Potential values Emes read on pH-meter for NaCl solutions of Im = 0.01 m to 3.19
	m at pcH = 2 using the combined-glass electrode with the filling solution of saturated KCl:
	liquid junction potentials ∆Emes calculated to estimate the potential differences between the solution of Im = 3.19 m and those of different Im; ∆Emes then used in the four-point calibration HCl 0.01M + NaCl 0.11M 1.0020 b 0.12 275.4 24.1
	with commercial buffer solution of pH 1.64, 4.01, 6.87, 9.18, assumed as diluted solutions of Im
	≈ 0.01 m. HCl 0.01M	1.0000		0.01	275.0	24.5
	Solution	Conversion	factor	Im (mol kgw -1 ) Emes	∆Emes = Emes (Im = 3.19 m)
		(from molarity M to		(mV)	-Emes
		molality m)
	HCl 0.01 M + NaCl 2.99 M	1.0668 a		3.19	299.5	0.0
	HCl 0.01 M + NaCl 1.49 M	1.0319 a		1.56	286	13.5
	HCl 0.01 M + NaCl 0.99 M	1.0215 a		1.04	280.8	18.7
	HCl 0.01M + NaCl 0.74 M	1.0165 a		0.76	279.7	19.8
	HCl 0.01 M + NaCl 0.59 M	1.0135 b		0.61	277.9	21.6
	HCl 0.01 M + NaCl 0.49 M	1.0118 a		0.51	277.8	21.7
	HCl 0.01 M + NaCl 0.41 M	1.0092 b		0.42	276.7	22.8
	HCl 0.01M + NaCl 0.29M	1.0064 b		0.30	276.5	23.0

Table IV -

 IV S2 Experimental sample information -pH values, calculated Ringböm coefficients α, [Mg 2+ ] (mol kgw-1 ) and deduced F0 and τ; the sample solutions giving the slope of 1 are set in italics.

	Im 6	7.71	1.14	2.85E-02	6.38E+07	10.16
	NaClO4 7	7.68	1.24	2.21 3.21E-02	7.95E+07	9.40
	(mol•kgw -1 ) 8	7.65	1.38	4.22E-02	7.53E+07	10.44
	Sample 9	pH 7.61	α 1.64	[Mg 2+ ] 5.67E-02	F0 (counts) τ (ns) 8.12E+07 10.20
	10 11	value 7.53 7.52	2.59 2.78	(mol•kgw -6.65E-02 1 ) 7.83E-02	5.81E+07 6.96E+07	12.89 12.48
	1 12	9.00 7.50	1.00 3.20	0 1.05E-01	3.86E+07 7.25E+07	9.27 12.65
	2	8.30	1.02	3.11E-03	3.19E+07	10.77
	3 4 Table IV -S3 Characteristic decomposition parameters of the luminescence spectrum in Fig. 8.23 1.05 1.53E-02 3.80E+07 10.40 8.16 1.06 4.31E-02 3.19E+07 12.31 5 7.85 1.02 5.86E-02 3.27E+07 13.60 IV -4 of the main text.
	6	7.70	1.07	6.50E-02	3.32E+07	15.74
	7 Peak Index	7.59	1.23 Area Fit	7.89E-02	3.47E+07 Center Max	16.70	Max Height		FWHM
	8 9 1 2 3 4 5 Im NaCl 6	7.54 7.51 7.49 7.46 7.43	1.40 1.57 8.94E+06 1.71 2.19E+06 1.98 2.02E+06 2.34 1.16E+06 4.68E+05 3.25E+05	9.80E-02 1.31E-01 1.59E-01 1.80E-01 2.11E-01 0.10	3.23E+07 3.71E+07 464.62 3.64E+07 484.37 4.38E+07 504.75 4.45E+07 525.76 547.63 569.17	19.15 20.27 27.86 25.32 28.66	6.43E+04 1.34E+05 1.21E+05 6.27E+04 2.06E+04 6.81E+03	0.25	13.06 15.50 15.74 17.29 21.32 44.75
	(mol•kgw -1 )
	Sample	pH	α	[Mg 2+ ]	F0 (counts) τ (ns)	pH	α	[Mg 2+ ]	F0 (counts) τ (ns)
		value		(mol•kgw -			value		(mol•kgw -
				1 )					1 )
	1	9.17	1.00	0	2.61E+07	9.16	9.00	1.00	0	4.82E+07	9.93
	2	8.26	1.09	3.48E-04	2.82E+07	13.15	8.30	1.05	8.23E-04	4.52E+07	9.72
	3	8.22	1.12	7.32E-04	2.37E+07	14.38	8.23	1.04	1.72E-03	2.49E+07	13.59
	4 Im NaClO4 5 6 (mol•kgw -1 ) Sample 7 8 9 1 2 3 4 Im 5 NaCl 6 (mol•kgw -1 ) 7 Sample 8 9 10 1 11 2 3 Im NaClO4 4 5 (mol•kgw -1 ) Sample 6 7 8 1 9 2 3 4 5 Im 6 NaCl 7 (mol•kgw -1 ) 8 Sample 9 10 11 1 12 2 13 3 4	8.07 8.05 8.00 pH value 7.90 7.87 7.78 9.00 7.74 8.45 8.30 8.23 8.20 8.09 8.06 pH 8.01 value 7.84 7.76 9.00 7.73 8.45 7.95 7.87 7.75 pH value 7.70 7.68 7.63 9.00 7.60 8.45 7.59 8.30 7.57 8.23 7.55 8.16 7.70 7.67 7.64 pH 7.61 value 7.58 7.54 9.00 7.51 8.45 7.48 8.30 8.23	1.49 1.64 2.08 α 3.80 4.80 8.90 1.00 12.13 1.02 1.06 1.12 1.17 1.45 1.62 α 3.71 6.17 10.78 1.00 14.18 1.01 1.09 1.19 1.70 α 2.20 2.45 3.35 1 4.17 1.01 4.36 1.02 5.05 1.05 5.89 1.06 1.27 1.38 1.54 α 1.74 2.02 2.53 1.00 3.05 1.01 3.73 1.03 1.03	1.25E-03 0.10 1.49E-03 2.38E-03 [Mg 2+ ] (mol•kgw -1 ) 4.16E-03 6.36E-03 1.56E-02 0 2.50E-02 1.76E-04 4.50E-04 1.57E-03 2.12E-03 0.50 4.45E-03 6.82E-03 [Mg 2+ ] 1.35E-02 (mol•kgw -2.36E-02 1 ) 4.18E-02 0 6.41E-02 8.50E-04 5.91E-03 1.05 6.34E-03 7.37E-03 [Mg 2+ ] (mol•kgw -1 ) 9.57E-03 1.50E-02 1.93E-02 0 2.32E-02 1.00E-03 3.07E-02 5.12E-03 4.07E-02 1.60E-02 6.08E-02 2.02E-02 2.68E-02 1.00 3.44E-02 4.50E-02 [Mg 2+ ] 6.30E-02 (mol•kgw -7.43E-02 1 ) 9.44E-02 0 1.20E-01 7.92E-04 1.46E-01 3.00E-03 1.18E-02	2.23E+07 2.49E+07 2.62E+07 F0 (counts) τ (ns) 18.58 19.11 20.35 2.90E+07 23.99 3.30E+07 25.72 3.51E+07 28.22 1.60E+07 8.59 9.00 8.00 7.95 7.88 pH value 7.81 7.79 7.77 4.59E+07 28.30 7.75 1.83E+07 13.28 8.23 7.73 2.29E+07 12.21 8.16 7.70 1.58E+07 16.71 8.09 1.59E+07 16.24 7.89 1.61E+07 23.44 7.82 2.07E+07 27.92 7.71 F0 (counts) τ (ns) pH 2.36E+07 28.39 7.66 value 2.43E+07 31.61 7.62 2.79E+07 33.48 7.58 7.96E+07 7.91 9.00 3.45E+07 34.83 7.53 8.48E+07 8.97 8.45 3.88E+07 11.65 8.30 2.97E+07 13.09 7.90 3.60E+07 13.09 7.82 F0 (counts) τ (ns) pH value 3.22E+07 14.37 7.73 3.25E+07 15.11 7.68 2.86E+07 17.26 7.65 3.33E+07 10.28 9.00 3.58E+07 16.70 7.62 2.49E+07 12.43 8.70 3.44E+07 17.94 7.59 2.57E+07 13.38 8.22 4.85E+07 19.69 7.57 2.57E+07 13.38 8.18 5.98E+07 18.11 7.55 1.88E+07 16.22 8.09 2.65E+07 15.07 7.98 2.36E+07 17.83 7.89 2.51E+07 18.62 7.78 F0 (counts) τ (ns) 2.75E+07 19.86 7.69 2.52E+07 22.79 7.64 3.18E+07 22.99 7.56 1.16E+08 7.44 3.27E+07 26.04 7.50 7.43E+07 8.58 3.54E+07 27.08 7.47 6.40E+07 8.84 6.22E+07 9.36	1.20 1.31 1.70 α 2.42 2.73 3.16 1.00 3.50 1.05 4.03 1.07 5.01 1.03 1.15 1.34 2.11 α 2.82 3.65 4.81 1.00 6.91 1.01 1.02 1.08 1.14 α 1.38 1.68 1.94 1.00 2.29 1.01 2.73 1.02 3.12 1.05 3.57 1.06 1.02 1.02 1.06 1.11 1.21 1.54 2.06 2.53	2.79E-03 3.66E-03 0.51 4.08E-03 [Mg 2+ ] 6.10E-03 8.46E-03 (mol•kgw -1.02E-02 1 ) 0 1.27E-02 9.46E-03 1.65E-02 1.93E-02 2.10E-02 2.14E-02 2.69E-02 0.75 3.61E-02 4.69E-02 [Mg 2+ ] 6.45E-02 (mol•kgw -1.13E-01 1 ) 1.16E-01 0 1.87E-01 8.19E-04 3.17E-03 6.39E-03 1.62 2.87E-02 [Mg 2+ ] 3.61E-02 5.03E-02 (mol•kgw -6.23E-02 1 ) 0 7.72E-02 9.38E-04 8.44E-02 2.37E-02 9.54E-02 5.52E-02 1.10E-01 1.43E-02 2.32E-02 2.60E-02 3.36E-02 4.55E-02 6.19E-02 8.60E-02 1.08E-01 1.39E-01	2.54E+07 2.76E+07 2.97E+07 3.23E+07 F0 (counts) τ (ns) 15.33 15.61 15.79 15.60 3.22E+07 18.31 3.42E+07 20.05 2.23E+07 3.85E+07 20.36 10.42 3.68E+07 3.83E+07 22.22 11.46 3.17E+07 4.05E+07 23.07 12.65 2.04E+07 14.74 3.16E+07 13.79 3.12E+07 15.75 2.08E+07 F0 (counts) τ (ns) 19.50 2.58E+07 20.56 2.61E+07 23.21 2.99E+07 5.03E+07 9.48 25.55 3.35E+07 6.70E+07 9.59 26.53 6.21E+07 9.63 5.82E+07 9.93 4.83E+07 11.22 5.51E+07 10.75 F0 (counts) τ (ns) 5.64E+07 11.11 5.74E+07 11.65 5.44E+07 5.83E+07 12.28 8.19 3.56E+07 5.23E+07 13.18 11.35 4.32E+07 5.12E+07 14.26 15.04 4.49E+07 6.07E+07 13.48 17.32 4.23E+07 14.30 4.20E+07 15.54 3.74E+07 17.42 4.13E+07 18.62 4.11E+07 20.13 4.67E+07 23.50 4.67E+07 22.17 5.05E+07 23.72 5.80E+07 24.88
	5	8.16	1.03	1.71E-02	4.17E+07	10.73

Table IV -

 IV S4 Formation constants log10K 0 used in Fig. IV -S1. et al.; 2 b Plummer and Busenberg; 3 c Palmer and Wesolowski; 4 d Shock, et al., 5 e Wagman, et al. 6 f Yeatts and Marshall; 7 ; g Turner, et al.; 8 h Olin, et al. 9

	log10K 0	Mg	Ca
	MCO3(aq)	2.98 a	3.22 b
	MOH +	2.32 c	1.17 d
	MF +	1.35 a	0.68 a
	MSO4(aq)	2.26 e	2.02 f
	MPO4 -	4.85 g	6.46 g
	MHPO4(aq)	2.88 g	2.55 g
	MH2PO4 +	1.17 g	1.40 g
	MSeO4(aq)	2.2 h	2 h
	a Sverjensky,

Table IV -

 IV S5 Evolutions of log10K° values from Mg 2+ to Ba 2+ with different ligands, used in Fig. IV -S2. Palmer and Wesolowski; 4 c Shock, et al.; 5 d Wagman, et al.; 6 e Plummer and Busenberg; 3 f Busenberg, et al.; 10 g Busenberg and Plummer; 11 h Prapaipong, et al.; 12 i Yeatts and Marshall; 7 j Reardon and Armstrong; 13 k Reardon and Armstrong; 13 l Davies; 14 m Dong and Brooks; 15 n Shang and Reiller 16

	log10K°1.1.3	Mg 2+	Ca 2+	Sr 2+	Ba 2+
	F -	1.35 a	0.68 a	0.14 a	-0.18 a
	OH -	2.32 b	1.17 c	0.69 c	0.49 d
	CO3 2-	2.98 a	3.22 e	2.81 f	2.71 g
	Ox 2-	3.43 h	3.19 h	2.92 h	2.93 h
	SO4 2-	2.26 d	2.02 i	2.30 j	2.70 k
	S2O3 2-	1.84 l	2.02 l	2.04 l	2.3 l
		4.27 m	5.34 m	5.02 m	4.84 m
	UO2(CO3)3 4-	4.73 p.w. in NaCl	5.36 in NaCl n
		4.43 p.w. in NaClO4	5.42 in NaClO4 n
	a Sverjensky, et al.; 2 b

Table IV -

 IV S6 Seawater composition from Millero, et al.17

	Temp	25°C
	pH	8.1
	pe	4
	Density (g/cm 3 ) 1.023
	Units	mol kgw -1
	Na	0.4861
	Mg	0.0547
	Ca	0.0106
	K	0.0106
	Sr	0.0001
	Cl	0.5658
	S(6)	0.0293
	Br	0.0018
	F	0.0001
	U	1.386E-08

Table IV -

 IV S7 Thermodynamic constants used in theoretical speciation of seawater case in Fig. IV -8 of the main text and in Fig. IV -S4.

Table IV -

 IV S9 Compositions of selected Indian drinking water samples, collected from the work of Coyte, et al.24

	Description	Units	Jai-112	Tonk-110	Bundi-104	Tonk-104	Jai-24
	Temperature	°C	29.6	29.7	27.1	29.1	30.5
	pH		7.4	7.15	7.68	7.55	6.83
	F	mg/L	1.83	2.68	5.15	5.25	0
	Cl	mg/L	293	314	388	1050	545
	Br	mg/L	0.58	1.6	1.18	2.11	1.15
	I	mg/L	0	0	1.512	0	0
	N(5) as NO3 -	mg/L	23.1	207	31.6	28.6	252
	S(6) as SO4 2-	mg/L	166	119	337	400	90.3
	C(4) as HCO3 -	mg/L	1530	737	970	1030	832
	Ca	mg/L	13.9	38.9	20.5	14.9	212
	Mg	mg/L	23	38	39.3	21.5	119
	Sr	mg/L	0.99	1.45	0.92	1.58	2.89
	Na	mg/L	800	484	700	1320	372
	Fe	mg/L	0	0	0	0	0.53
	Ba	mg/L	0.27	0.24	0.11	0.02	0.38
	Mn	mg/L	0	0	0	0	0.45
	Si	mg/L	10.5	26.8	12.4	6.31	23
	Li	μg/L	0	11.26	30.35	31.33	13.28
	B	μg/L	2310	754	2180	4140	893
	V	μg/L	0	0	0	0	0
	Cr	μg/L	0	0	0	0	0
	Ni	μg/L	0	0	0	0	3.79
	Cu	μg/L	0	0	0	0	20.69
	Zn	μg/L	392.15	10.8	13.1	22.13	0
	As	μg/L	0	0	0	0	0
	Se	μg/L	0	0	0	0	0
	Mo	μg/L	52.31	6.64	29.19	15.25	2.63
	Pb	μg/L	0.69	0.52	0.63	1.06	0.34
	U	μg/L	181.28	32.71	50.45	319.67	31.23

  used to recover the spectra. The analysed global data matrix at a specified temperature is composed of 1024 rows and n columns, with 1024 corresponding to the number of pixels of the CCD camera (416.46-633.81 nm wavelength range), and n corresponding to the sample number in each series. The calculations have been carried out with codes developed in-house

in R environment (https://github.com/ppernot/SK-Ana). Unlike the routine application of the developed codes to spectro-kinetics matrices databuilt by delay-varying measurements -, the global matrices in this work were constructed from the spectra of different [Ca 2+ ] concentrations, measured at D = 25 ns. By this means, the problem can be recast into a decomposable form that emphasizes the bilinearity between Ca 2+ -dependent concentration and spectral properties, as expressed in the following equation:

Table V -

 V 1 Formation constants log10Kn.1.3 derived from the rounded off slopes in Fig. V -6 and Fig. V -7 with values of log10K0.1.3° taken from Götz et al.49

	CanUO2(CO3)3 (4-2n)-
	T (°C)	log10K1.1.3	log10K2.1.3	log10K°1.1.3	log10K°2.1.3 log10K°0.1.3 log10β°1.1.3	log10β°2.1.3
	10	3.92 ± 0.03	6.73 ± 0.03	5.62 ± 0.03	9.31± 0.03 22.28 ± 0.04 27.91 ± 0.04 31.59± 0.04
	15	3.85 ± 0.03	6.58 ± 0.05	5.57 ± 0.03	9.17± 0.05 22.11 ± 0.04 27.68 ± 0.04 31.28± 0.05
	22 34	3.56 ± 0.02	5.93 ± 0.03	5.29 ± 0.02	8.55± 0.03 21.93 ± 0.04 27.22 ± 0.04 30.48± 0.05
	25	3.69 ± 0.03	6.23 ± 0.03	5.43 ± 0.03	8.87± 0.03 21.81 ± 0.04 27.24 ± 0.04 30.68± 0.04
	35	3.75 ± 0.05	6.52 ± 0.05	5.52 ± 0.05	9.18± 0.05 21.58 ± 0.04 27.10 ± 0.05 30.76 ± 0.04
	40	3.89 ± 0.02	6.41 ± 0.02	5.68 ± 0.02	9.10± 0.02 21.48 ± 0.04 27.16 ± 0.04 30.58± 0.04
	45	3.78 ± 0.01	6.46 ± 0.04	5.59 ± 0.01	9.17± 0.04 21.40 ± 0.04 26.99 ± 0.04 30.57± 0.04
	50	3.80 ± 0.03	6.23 ± 0.04	5.63 ± 0.03	8.97± 0.04 21.33 ± 0.04 26.96 ± 0.04 30.30± 0.04
	MgUO2(CO3)3 2-
	T (°C)	log10K1.1.3	log10K°1.1.3	log10K°0.1.3	log10β°1.1.3
	5	3.00 ± 0.03	4.67 ± 0.03	22.48 ± 0.04 27.18 ± 0.04
	10 15	3.04 ± 0.07 2.76 ± 0.03	4.72 ± 0.07 4.45 ± 0.03	18 22.28 ± 0.04 27.03 ± 0.07 15551 22.11 ± 0.04 26.59 ± 0.04
	22 36	2.98 ± 0.02	4.68 ± 0.02	21.93 ± 0.04 26.65 ± 0.04
	30	2.73 ± 0.05	4.46 ± 0.05	21.69 ± 0.04 26.32 ± 0.05

Table V -

 V 2).kJ mol-1 , and for Ca2UO2(CO3)3(aq) with ∆r CHC H°m = -(10.992 ± 14.150) kJ mol -1 and ∆r Exo H°m = -(16.209 ± 9.064) kJ mol -1 . We specifically calculated the thermodynamic constants for the successive complexation of Ca 2+ to CaUO2(CO3)3 2-by subtracting the values determined for CaUO2(CO3)3 2-from those for Ca2UO2(CO3)3(aq) (TableV -2). The uncertainties were apparently too large, due to the very weak variation of the values of log10K°n.1.3 in the investigated temperature ranges.One can note that for the formation of CaUO2(CO3)3 2-, the entropy term T∆rS°m varies from (27.79 ± 5.40) kJ mol -1 to (36.45 ± 5.96) kJ mol -1 at room temperature depending on the fitting equation. It seems though that the formation of CaUO2(CO3)3 2-is entropy-driven, consistent with the trends observed both by Endrizzi and Rao18 and Jo et al.28 , whilst the enthalpy value is closely related to the approximation approach. In comparison, the second complexation of CaUO2(CO3)3 2-by Ca 2+ is always exothermic regardless of the approximation approach. The entropy term T∆rS°m of this reactionfrom (9.60 ± 10.56) kJ mol -1 to (13.80 ± 13.30) kJ mol - 1shares nearly equal importance with the enthalpy term but is much less important than in the first complexation.

Table V

 V

							-1; the values of
	log10K°n.1.3 at 70°C calculated by PhreeqC and used to perfom the uranium speciation, shown
	in Fig. V -10.
	Isoelectric	A1	A2	A3	A4	A5	log10K°n.1.3	at
	reaction						70°C
	CaUO2(CO3)3 2-	33822.8	6.01794	-1647050	-12537.7759 83260000	5.81
	Ca2UO2(CO3)3(aq	35198.8	6.25537	-1722880	-13040.2035 88090000	9.42
	) MgUO2(CO3)3 2-	33994.3	6.04948	-1654900	-12602.4447 83660000	4.69

Table V -

 V S4 Main specific ion interaction coefficients used in this work.

	T (°C)	30
	Specific ion interaction coefficient Sample pH α [Mg 2+ ]	Value (± 1σ) F0		Ref.
	ε(H + , Cl -)	value		(mol•kgw -1 )	0.12 ± 0.01 (counts)		Guillaumont, et al. 2
	ε(Na + , CO3 2-)				-0.08 ± 0.03		Guillaumont, et al. 2
	ε(UO2 2+ , Cl -)	8.97	1.00	1.00E-05	0.46 ± 0.03 9.65E+07		Guillaumont, et al. 2
	ε(Mg 2+ , Cl -) ε(UO2(CO3)2 2-, Na + ) 8.80 8.42 ε(UO2(CO3)3 4-, Na + ) 8.38 8.25	1.00 1.03 1.04 1.10	1.57E-04 2.37E-04 5.31E-04 1.18E-03	0.19 ± 0.02 4.32E+07 -0.02 ± 0.09 8.18E+07 -0.01 ± 0.11 3.85E+07 5.23E+07		Guillaumont, et al. 2 Guillaumont, et al. 2 Guillaumont, et al. 2
	ε(MgUO2(CO3)3 2-, Na + ) (in NaCl) 8.09 1.55 1.18E-03	0.31 ± 0.02 5.59E+07		Shang et al. 4
	ε(CaUO2(CO3)3 2-, Na + ) (in NaCl) 8.05 1.84 2.21E-03 ε(Ca2UO2(CO3)3, NaCl) 7.90 4.49 3.38E-03 7.81 8.56 4.58E-03	0.29 ± 0.11 4.96E+07 0.66 ± 0.22 4.24E+07 6.01E+07		Shang et al. 5 Shang et al. 5
		7.78	10.72	1.34E-02	1.04E+08
	Table V -S5 (a) Experimental sample information: pH values, calculated Ringböm coefficients 7.75 13.48 1.51E-02 1.23E+08
	α -see Shang et al. 4-6 for details -, [Mg 2+ ] (mol kgw -1 ) and deduced F0; the sample solutions 7.72 16.98 3.02E-02 1.95E+08
	giving the slope of 1 are set in italics.
	T (°C)	5				10
	Sample	pH	α	[Mg 2+ ]	F0	pH	α	[Mg 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
	1	8.92	1.00	0	1.46E+08	8.99	1.00	1.00E-05	1.26E+08
	2	8.59	1.00	1.57E-04	1.16E+08	8.61	1.00	1.57E-04	1.45E+08
	3	8.29	1.01	2.37E-04	1.23E+08	8.30	1.01	2.37E-04	1.29E+08
	4	8.24	1.02	5.31E-04	1.16E+08	8.26	1.01	5.31E-04	1.66E+08
	5	8.13	1.08	7.50E-04	1.11E+08	8.18	1.03	1.18E-03	1.35E+08
	6	7.99	1.48	1.18E-03	1.13E+08	8.05	1.19	1.18E-03	1.24E+08
	7	7.88	2.57	1.82E-03	6.41E+07	7.89	1.31	1.82E-03	1.06E+08
	8	7.82	3.78	2.21E-03	1.11E+08	7.82	2.63	2.21E-03	1.42E+08
	9	7.79	4.65	3.38E-03	1.23E+08	7.79	5.13	3.38E-03	1.04E+08
	10	7.72	7.74	4.58E-03	1.14E+08	7.72	7.96	4.58E-03	1.16E+08
	11	7.69	9.70	8.89E-03	1.37E+08	7.69	12.55	8.89E-03	1.77E+08
	12	7.64	14.25	1.34E-02	1.32E+08	7.64	16.95	1.34E-02	2.06E+08
	13	7.61	17.99	1.51E-02	1.28E+08	7.61	25.30	1.51E-02	2.39E+08
	14	7.54	31.21	3.02E-02	1.29E+08	7.54	34.37	3.02E-02	2.26E+08
	T (°C)	15				22 4
	Sample	pH	α	[Mg 2+ ]	F0	pH	α	[Mg 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
	1	8.97	1.00	1.00E-05	7.97E+07	9.17	1.00	0	2.61E+07
	2	8.80	1.00	1.57E-04	6.20E+07	8.26	1.09	3.48E-04	2.82E+07
	3	8.42	1.01	2.37E-04	6.38E+07	8.22	1.12	7.32E-04	2.37E+07
	4	8.38	1.01	5.31E-04	5.52E+07	8.07	1.49	1.25E-03	2.23E+07
	5	8.26	1.04	1.18E-03	5.04E+07	8.05	1.64	1.49E-03	2.49E+07
	6	8.10	1.22	1.18E-03	5.38E+07	8.00	2.08	2.38E-03	2.62E+07
	7	8.05	1.41	1.82E-03	5.74E+07	7.90	3.80	4.16E-03	2.90E+07
	8	7.90	3.00	2.21E-03	6.08E+07	7.87	4.80	6.36E-03	3.30E+07
	9	7.86	3.90	3.38E-03	5.95E+07	7.78	8.90	1.56E-02	3.51E+07
	10	7.78	6.88	4.58E-03	6.21E+07	7.74	12.13 2.50E-02	4.59E+07
	11	7.72	10.80	8.89E-03	7.20E+07

Table V -

 V S5 (b). Experimental sample information -pH values, calculated Ringböm coefficients α, [Ca 2+ ] (mol kgw-1 ) and deduced F0; the sample solutions giving the slope of 1 are set in italics; the sample solutions giving the slope of 2 are set in bold.

	T (°C)	10				15
	Sample	pH	α	[Ca 2+ ]	F0	pH	α	[Ca 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
		8.97	1.00	0	3.42E+08 8.97	1.00	0	5.22E+08
		8.50	1.00	5.72E-05	2.92E+08 8.50	1.01	5.72E-05	3.75E+08
		8.48	1.00	6.18E-05	3.10E+08 8.48	1.01	6.18E-05	3.65E+08
		8.36	1.01	6.48E-05	3.15E+08 8.36	1.02	6.46E-05	3.91E+08
		8.30	1.01	8.06E-05	3.42E+08 8.20	1.07	8.72E-05	4.21E+08
		8.22	1.03	9.87E-05	3.50E+08 8.16	1.11	1.19E-04	4.43E+08
		8.18	1.05	1.23E-04	3.66E+08 8.12	1.18	1.56E-04	4.44E+08
		8.08	1.18	1.57E-04	3.80E+08 8.05	1.41	1.74E-04	4.92E+08
		7.99	1.55	2.31E-04	4.29E+08 8.00	1.72	2.45E-04	5.19E+08
		7.92	2.17	4.80E-04	4.51E+08 7.93	2.49	4.34E-04	5.30E+08
		7.82	4.08	7.41E-04	5.25E+08 7.87	3.64	6.65E-04	5.35E+08
		7

.78 5.43 1.12E-03 5.28E+08 7.83 4.80 1.15E-03 5.42E+08 7.75 6.75 1.33E-03 5.68E+08 7.76 7.98 1.76E-03 5.78E+08 7.70 9.82 2.00E-03 6.02E+08 7.70 12.59 1.96E-03 5.94E+08 7.62 18.17 2.63E-03 6.13E+08 7.62 23.45 2.26E-03 6.58E+08 7.56 29.13 3.34E-03 6.84E+08 7.57 34.79 3.86E-03

		7.81	7.07	3.42E-03	3.08E+08
		7.75	11.55	4.70E-03	3.53E+08
		7.63	28.25	6.25E-03	3.43E+08
		7.60	37.08	8.55E-03	3.74E+08
		7.55	53.34	1.04E-02	3.84E+08
		7.52	70.02	1.53E-02	4.25E+08
		7.50	79.91	1.74E-02	4.30E+08
		7.48	96.30	2.49E-02	5.85E+08
	T (°C)	35				40
	Sample	pH	α	[Ca 2+ ]	F0	pH	α	[Ca 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
		8.97	1.00	0	3.86E+08 8.97	1.00	0	2.85E+07
		8.58	1.01	2.69E-05	1.71E+08 8.36	1.08	1.94E-04	5.43E+07
		8.55	1.02	4.67E-05	1.04E+08 8.16	1.42	2.98E-04	5.78E+07
		8.50	1.02	5.72E-05	3.66E+08 8.12	1.62	4.19E-04	5.89E+07
		8.48	1.03	6.18E-05	3.49E+08 8.05	2.23	6.13E-04	6.24E+07
		8.36	1.06	6.48E-05	2.07E+08 8.00	2.96	1.01E-03	6.27E+07
		8.16	1.33	1.39E-04	2.90E+08 7.93	4.66	1.80E-03	7.83E+07
		8.12	1.50	1.96E-04	2.98E+08 7.87	7.11	2.56E-03	7.96E+07
		8.05	2.02	2.87E-04	3.20E+08 7.81	11.09	3.30E-03	8.15E+07
		8.00	2.65	4.71E-04	3.03E+08 7.78	13.93	4.25E-03	9.20E+07
		7.93	4.13	8.44E-04	3.19E+08 7.75	17.53	5.59E-03	1.29E+08
		7.87	6.28	1.09E-03	3.09E+08 7.71	23.90	7.04E-03	1.43E+08
		7.78	12.24	1.80E-03	3.11E+08 7.68	30.21	7.77E-03	1.58E+08
		7.70	22.69	2.15E-03	3.73E+08
		7.62	42.54	3.25E-03	3.37E+08
		7.57	80.28	4.50E-03	3.41E+08	8.48
		7.49	119.6	7.20E-03	3.79E+08
	T (°C)	45	4			50
	Sample	pH	α	[Ca 2+ ]	F0	pH	α	[Ca 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
		8.97	1.00	0	2.64E+07 8.97	1.00	0	2.76E+07
		8.48	1.05	1.10E-04	3.20E+07 8.16	1.64	2.98E-04	2.79E+07
		8.36	1.10	1.94E-04	4.81E+07 8.12	1.92	4.19E-04	3.01E+07
		8.16	1.52	2.98E-04	3.08E+07 7.98	4.20	6.13E-04	3.02E+07
		8.12	1.76	4.19E-04	3.82E+07 7.95	5.14	1.01E-03	3.99E+07
		8.05	2.47	6.13E-04	4.30E+07 7.92	6.33	1.80E-03	5.58E+07
		8.00	3.31	1.01E-03	4.34E+07 7.88	8.42	2.56E-03	5.14E+07
		7.93	5.24	1.80E-03	5.57E+07 7
									6.74E+08
		7.51	43.34	4.40E-03	7.77E+08 7.50	60.70	4.28E-03	7.12E+08
	T (°C)	22 5				25
	Sample	pH	α	[Ca 2+ ]	F0	pH	α	[Ca 2+ ]	F0 (counts)
		value		(mol•kgw -1 )	(counts)	value		(mol•kgw -1 )
		9.90	1.00	0	1.70E+08 8.97	1.00	0	3.77E+08
		8.66	1.06	2.78E-05	1.66E+08 8.36	1.03	6.48E-05	3.04E+08
		8.42	1.04	9.09E-05	1.29E+08 8.16	1.19	1.39E-04	3.27E+08
		8.26	1.10	1.76E-04	1.75E+08 8.12	1.30	1.96E-04	3.44E+08
		8.23	1.12	2.22E-04	2.25E+08 8.05	1.65	2.87E-04	3.70E+08
		8.15	1.24	3.06E-04	2.40E+08 8.00	2.10	4.71E-04	4.10E+08
		8.11	1.37	4.04E-04	2.27E+08 7.93	3.19	8.44E-04	3.91E+08
		8.10	1.34	4.93E-04	2.18E+08 7.87	4.76	1.09E-03	3.84E+08
		8.08	1.51	6.00E-04	2.74E+08 7.83	6.33	1.87E-03	3.69E+08
		8.00	2.06	9.03E-04	2.47E+08 7.76	10.64	2.69E-03	3.88E+08
		7.97	2.52	1.22E-03	3.08E+08 7.70	16.88	3.02E-03	4.00E+08
		7.92	3.30	1.72E-03	3.12E+08 7.62	31.61	4.07E-03	4.19E+08
		7.87	4.79	2.08E-03	3.32E+08 7.57	46.99	5.72E-03	3.92E+08
		7.85	5.51	2.64E-03	3.55E+08 7.50	82.02	8.23E-03	4.33E+08

.85 10.50 3.30E-03 5.58E+07 7.82 8.04 2.56E-03 5.64E+07 7.80 15.29 4.25E-03 6.60E+07 7.78 15.78 3.30E-03 7.19E+07 7.75 22.42 5.59E-03 7.22E+07 7.72 19.87 4.25E-03 7.94E+07 7.73 26.18 7.70E-03 8.28E+07 7.69 27.09 5.59E-03 1.00E+08 7.70 33.06 9.09E-03

				9.48E+07
	7.65	34.25	7.70E-03	1.00E+08
	7.62	43.35	9.09E-03	1.20E+08

Table V -

 V S6 Values of the Debye-Hückel parameters A(T) at different temperatures, taken fromthe NEA-TDB reviews.7

	t/°C (p = 1bar) A(T) / kg 1/2 mol -1/2
	5	0.494
	10	0.497
	15	0.501
	20	0.505
	25	0.509
	30	0.513
	35	0.518
	40	0.523
	50	0.533
	75	0.564

Table V -

 V S7 The formation constants of log10K013°, log10Q1 and log10Q2 in the temperature range of 5-90°C, calculated using the exiting general formula; the global formation constants of log10β113°(Ca), log10β213°(Ca) and log10β°113 (Mg) were estimated with the Van't Hoff Equation; the formation constants of log10K°113 (Ca), log10K°213(Ca), log10K°113 (Mg) were finally determined by subtracting the sum of log10K°013, log10Q1 and log10Q2 from the global equilibrium constants log10β°113 (Ca), log10β°213 (Ca) and log10β°113 (Mg).

	T	log10K°013	log10Q1	log10Q2	log10β°113(Ca)	log10β°213(Ca)
	(°C)	Götz, et al. 8	Patterson, et al. 9	Patterson, et
				al. 10
	5	22.48	-6.51	-10.56	-6.06	-19.35
	10	22.28	-6.46	-10.49	-6.04	-19.29
	15	22.11	-6.41	-10.43	-6.03	-19.23
	20	21.95	-6.38	-10.38	-6.02	-19.18
	22	21.93	-6.37	-10.36	-6.02	-19.16
	25	21.81	-6.35	-10.34	-6.01	-19.13
	35	21.58	-6.31	-10.26	-5.99	-19.08
	40	21.48	-6.29	-10.23	-5.98	-18.98
	45	21.40	-6.28	-10.20	-5.98	-18.94
	50	21.33	-6.28	-10.18	-5.97	-18.89
	55	21.27	-6.28	-10.16	-5.96	-18.85
	60	21.22	-6.28	-10.15	-5.95	-18.80
	65	21.18	-6.29	-10.13	-5.94	-18.76
	70	21.15	-6.29	-10.12	-5.94	-18.72
	75	20.96	-6.30	-10.12	-5.93	-18.69
	80	20.88	-6.32	-10.11	-5.92	-18.65
	85	20.81	-6.33	-10.11	-5.91	-18.61
	90	20.73	-6.35	-10.11	-5.91	-18.58
	T	log10β°113(Mg) log10K°113(Ca)	log10K°213(Ca)	log10K°113°(Mg)
	(°C)
	5	-6.93	5.60	9.37	4.72
	10	-6.95	5.57	9.27	4.67
	15	-6.96	5.55	9.20	4.63
	20	-6.97	5.55	9.15	4.60
	22	-6.97	5.51	9.10	4.55
	25	-6.98	5.55	9.12	4.58
	35	-6.99	5.56	9.18	4.57
	40	-7.00	5.58	9.11	4.56
	45	-7.01	5.60	9.13	4.55

Table V -

 V S8 Characteristic decomposition parameters of the pure spectra of CanUO2(CO3)3(4- 2n)-and MgUO2(CO3)32-shown in Fig. V -S2 and Fig. V -S3.

	Peak index	Peak Type	FWHM	Center	Area IntgP
	CaUO2(CO3)3 2-
		Gaussian	23.10	479.58	14.98
		Gaussian	17.97241	496.618	28.22
		Gaussian	17.17688	514.51	24.26
		Gaussian	23.7694	533.45	16.11
		Gaussian	52.32048	552.16	16.42
	Peak index	Peak Type	FWHM	Center	Area IntgP
	Ca2UO2(CO3)3(aq)
		Gaussian	13.65	465.01	14.54
		Gaussian	13.69	484.14	27.03
		Gaussian	15.97	504.62	28.42
		Gaussian	17.37	525.98	17.13
		Gaussian	23.51	548.13	9.13
		Gaussian	34.64	576.52	3.76
	Peak index	Peak Type	FWHM	Center	Area IntgP
	MgUO2(CO3)3 2-
		Gaussian	24.15	465.60	21.44
		Gaussian	15.45	484.98	21.17
		Gaussian	17.89	502.85	22.221
		Gaussian	26.34	523	16.81
		Gaussian	59.05	545.93	18.38

  Table VI -1, are considerably variable and reflectthe nature and effects of geochemical processes on groundwater composition. Groundwaters usually have much higher concentrations of principal constituents than surface waters. Those in contact with rock tend to be more concentrated than shallow waters.

* http://weppi.gtk.fi/publ/foregsatlas/

Table VI -

 VI 3 Percentages of major uranium carbonate species and the activity of important ions

	Samples	"untreated" water	acidified water at pH 6.5
	CaUO2(CO3)3 2-	50.6%	41.6%
	MgUO2(CO3)3 2-	26.1%	6.7%
	Ca2UO2(CO3)3(aq)	19.7%	50.8%
	UO2(CO3)3 4-	3.5%	0.9%
	Total	99.89%	99.97%
	lg P(CO2(g))	-2.92	-2.73
	lg a(Ca 2+ )	-3.446	-2.946
	lg a(Mg 2+ )	-2.934	-2.941
	lg a(HCO3 -)	-2.44	-2.386
	lg a(Na + )	-2.15	-2.164
	lg a(H + )	-8.3	-8.16
	lg a(Ca 2+ ) -2 lg a(H + )	13.15	13.37
	lg a(Mg 2+ ) -2 lg a(H + )	13.67	13.38
	n = a(Ca 2+ )/(a(Ca 2+ ) + a(Mg 2+ ))	0.76	0.50
	nlg a(Mg 2+ ) + (1-n) lg a(Ca 2+ ) -2 lg
	a(H + )	13.55	13.38
	lg a(HCO3 -) + lg a(H + )	-10.74	-10.55
	lg a(Na + ) -lg a(H + )	6.15	6.00
	I (mol kgw -1 )	0.022	0.027
	6.2 TRLFS measurement

-3 816.130 ± 3.050 p.w. -3 816.015 ± 2.991 p.w. multilinear regression

Remerciements

Acknowledgment

This work was financed by ONDRAF-NIRAS (contract DEN4857-CCHO 2018-0456/00/00). Dr. Hélène Isnard is acknowledged for her help during the dissolution of U3O8.

Acknowledgments

This work was financed by CEA and ONDRAF-NIRAS (contract DEN4857-CCHO 2018-0456/00/00). Dr. Hélène Isnard is acknowledged for her help in the dissolution of U3O8.

III -6 References

1.

G. Bernhard, G. Geipel, V. Brendler and H. Nitsche, Radiochim. Acta, 1996, 74, 87-91.

Acknowledgment

This work was financed by CEA and ONDRAF-NIRAS (contract DEN4857-CCHO 2018-0456/00/00). Dr. Hélène Isnard is acknowledged for her help in the dissolution of U3O8.

Acknowledgements

The authors thanks Dr. Pascal Pernot, who developed the MCR-ALS analysis code in the R environment, for his help; and Dr. Hélène Isnard who prepared the uranium(VI) solution.

Mettre le remerciement pour Stéphane Brassinnes (ONDRAF

Funding sources

This work was financed by ONDRAF-NIRAS and CEA (contract DEN4857-CCHO 2018-0456/00/00).

 4-+ 2H + ⇄ UO2(CO3)2 2-+ CO2(g) + H2O

12.92 ± 0.15* Recalculated from 2UO2(CO3)3 4-+ 7H + ⇄ (UO2)2CO3(OH)3 -+ 5CO2(g) + 2H2O 46.21 ± 0.70* Recalculated from Ca 2+ + UO2(CO3)3 4-⇄ CaUO2(CO3)3 2- 5.36 ± 0.01 a 8 in NaCl 2Ca 2+ + UO2(CO3)3 4 8 The correct value was reported in the text.

Table III -2 Main specific ion interaction coefficients used in this work.

Specific ion interaction coefficient Value (± 1σ) kgw mol -1 Ref. ε(H + ,ClO4 -) 0.14 ± 0.02 29 ε(Na + ,CO3 3,NaClO4) 0.18 ± 0.07 this work individual regression ε(CaUO2(CO3)3 2-,Na + ) (in NaCl) 0.32 ± 0.12 this work multilinear regression ε(Ca2UO2(CO3) 3,NaCl) 0.66 ± 0.15 this work multilinear regression ε(CaUO2(CO3)3 2-,Na + ) (in NaClO4) 0.00 ± 0.10 this work multilinear regression ε(Ca2UO2(CO3) 3,NaClO4) 0.17 ± 0.12 this work multilinear regression Following the expression of Kn.1.3, the fluorescence intensity of UO2(CO3)3 4-should be measured at the same pH as in the sample of CanUO2(CO3)3 (4-2n)-. However, the fluorescence spectrum of a free-calcium triscarbonatouranyl sample under weakly basic conditions (pH < 8)

is interfering inevitably with hydroxo, less-complexed carbonatouranyl, and polynuclear speciese.g., UO2(CO3)2 2-, UO2CO3(aq), UO2(OH) + , UO2(OH)2(aq), (UO2)2CO3(OH)3 -. As an thus cover the data better than one-parameter model, which generates a linear trend. Therefore, the goodness-of-fit, estimated by the sum of squares of distances between the data points and the best fit curvature, is naturally and technically better in more complicated models than that in the simpler model. 87 To discriminate one model from another among the current two-parameter models, the value of log10 1.1.3 determined at Im = 3.50 m is added in the figure to help predicting and choosing the best trend line. It appears that the model proposed by

Ciavatta better describes the data. 73 Nonetheless, there is a possibility that the value of log10 1.1.3 has been somewhat underestimated, because of the unclear quenching mechanisms at this ionic strength. This is also the reason why we did not take this value into the evaluations of log10 ° and ε. It seems now possible to state that the model constructed by Bretti, et al. 84,85 could serve as a trade-off between the one-parameter SIT model and that of Ciavatta. 73

III -5 Conclusion

The stability constants log10 ° and specific ion interaction coefficients ε were evaluated for where FI is the fluorescence intensity at a specific delay time and FI0 is at D = 0. The gate width W = 1 μs and τi is the decay-time of species i.

Correction for slope analysis and calculation of Ringböm coefficients

The introduction of Ringböm 49 coefficient α has been discussed in our previous studies. 24,25 We will simply recall here the equations used for the slope analysis and the meanings of α for explanatory purposes. The successive binding of Mg 2+ with UO2(CO3)3 4-is described by Eq. IV -2,

where the unit of concentration is molality (mol kgw -1 ).

The slope analysis, as expressed in Eq. IV -3, is used to evaluate the conditional successive formation constants log10Kn.1.3 after rounding-off the slope n, representative of stoichiometric number of Mg 2+ , to its nearest integer,

where R represents the concentration ratio between the [MgnUO2(CO3)3 (4-2n)-] and [UO2(CO3)3 4-] taking α into account.

The α coefficient allows calculating the amount of UO2(CO3)3 23 ; c Shang, et al. 24

Specific ion interaction Theory (SIT)

The formation constants log10K°n.1.3 were obtained by extrapolating the log10Kn.1.3 values at investigated ionic strengths with the SIT approach, in which the activity coefficient is expressed as follows,

where ε(j,k) (kgw mol -1 ) is the specific ion interaction coefficient between the aqueous ion j and the oppositely charge ion k in medium solution of ionic strength Im. The Debye-Hückel term is expressed in Eq. IV -5,

where the value of 0.507 kg 1/2 mol -1/2 was assigned to Debye-Hückel constant A, corresponding to the actual temperature of 22°C. 22 In Eq. IV -6, the relationship between the global conditional formation constants log10 β0. The specific ion interaction coefficients used in calculations are summarized in Table IV -3. It is noteworthy that ε(UO2 2+ ,Cl -) = 0.21 kgw mol -1 was not used in this study, because Guillaumont, et al. 22 proposed that the complexation between uranyl and chloride ion should not be ignored, though the interaction is rather weak. For the sake of consistency with our previous works, ε(UO2 2+ , ClO4 -) = 0.46 kgw mol -1 was used in this work.

The simultaneous fitting with a common value of log10β°1. The earlier published stability constants of MgUO2(CO3)3 2-complex are listed in Table IV -6.

The obtained constants in the present work are slightly higher than those reported previously in the case of NaCl. It is worth noting that the value of log10β°1.1.3 depends on the extrapolation models used to describe the correlation between activity coefficient and ionic strength. For the previous propositions, the mainly used models were Davies, 35 applicable to Im < 0.3 mol kgw -1 , and SIT, 22 valid for a wider range of Im < 3-4 mol kgw -1 . In a systematic manner, Dong and

Brooks 15 compared the values of log10β°1.1.3 using these two different models in NaNO3. The authors mentioned that the use of Davies equation is not appropriate in their study in which the maximum ionic strength extended to 0.509 mol kgw -1 NaNO3 and also proposed the results for ε(MgUO2(CO3)3 2-, Na + ) estimated by SIT as -3.0 ± 0.3 kgw mol -1 .

positions obtained by Gaussian-based decomposition of the spectra acquired with 300 lines mm -1 grating. A very slight bathochromic shift in the luminescence spectra when varying the principal aqueous component from Ca-UO2-CO3 to Mg-UO2-CO3 was already reported. 9,10,47 In this example, the properties of MnUO2(CO3)3 (4-2n)-, especially their formation constants and interaction coefficients, are essential to have reliable appreciation of U(VI) equilibrium speciation that is driven by topography after the chemical weathering at the site. The CanUO2(CO3)3 (4-2n)-species are proven to be responsible for the enhancement of U(VI) mobility by promoting its desorption from the clay minerals in literature. [68][69][70][71][72][73][74] Some emphasis should also be placed on Mg-concentrated natural conditions where the MgUO2(CO3)3 2-complexsimilar to CanUO2(CO3)3 (4-2n)-in terms of thermodynamic properties but did not receive sufficient attentioncould take a large proportion of aqueous U(VI).

Case of Indian drinking water

Groundwater in India, including the water directly provided to consumers without treatment, suffers from numerous quality issues. 41,42,75 The prevailing uranium in groundwater could represent a serious concern to the local people, in addition to multiple high concentrations of inorganic contaminants such as arsenic and halides. Coyte, et al. 41 compiled uranium data from several states in North-Western India where the groundwater resources have been severely

affected by over-exploitation and climate effects. In the shallower layers, the oxidizing and carbonate-rich conditions favour the formation of soluble uranyl carbonate complexes, which lead to high aqueous U(VI) concentration. The natural migration into deeper parts explains the observed considerable decrease in U(VI) with depth.

In the work of Coyte, et al., 41 speciation modelling were done for some wells using the llnl database provided with the PhreeqC package, 43,44 where the log10β°n.1.3 for Mg-, Ca-, Sr-uranyl carbonate complexes, taken from Dong and Brooks, 14 were manually added. The calculation

IV -6 Conclusion

The complexation behaviour of MgUO2(CO3)3 2-complex was studied in NaCl and NaClO4 media by collecting the luminescence spectra and determining their thermodynamic constants log10β°1.1.3 and specific ion interaction theory coefficients ε using TRLS. The luminescence spectra of MgUO2(CO3)3 2-complex showed slight hypochromic shifts compared with CaUO2(CO3)3 2-, detectable using the 1800 lines mm -1 grating with our apparatus. The equilibrium constants measured log10β°(MgUO2(CO3)3 2-) = 26.40 ± 0.07 and the specific ion interaction coefficients represent similar features as CanUO2(CO3)3 (4-2n)-complexes, i.e.

different values in NaClO4 -ε(MgUO2(CO3)3 2-, Na + ) = 0.19 ± 0.11 kgw mol -1and in NaCl -

The underlying mechanisms in the ion interactions between M 2+ (M = Mg, Ca) and UO2(CO3)3 4-were discussed and analysed from different aspects. The stable octahedral Mg(H2O)6 2+ structure may result in steric hindrance when coordinating to UO2(CO3)3 4-compared with the more hydrated Ca(H2O)7 2+ , which facilitates the water exchange. This intrinsic characteristic of magnesium also gives reason for the anomaly in the stability constants for the MUO2(CO3)3 2-complexes (M = Mg, Ca, Sr, Ba) observed otherwise.

Concerning the interaction coefficients, our results suggest that the ion associations occurring among Na + , M 2+ , Cl -and ClO4 -are supposed to be important at high Im. This work is providing a complete thermodynamic constants set for the MnUO2(CO3)3 (4-2n)-complexes, both in NaCl and NaClO4 systems.

The applications to several representative geochemical contexts highlight the necessity to account for the MgUO2(CO3)3 2-complex in theoretical speciation calculation. Particularly when the ratio of magnesium over calcium approaches or overpasses the ratio of the formation constants.

Chapter V. Effect of Temperature on the Complexation of Triscarbonatouranyl(VI) with Calcium and Magnesium in NaCl

Aqueous Solution.

V -1 Abstract

The complex formation of triscarbonatouranyl(VI) UO2(CO3)3 4-with alkaline earth metal ions Mg 2+ and Ca 2+ in 0.10 mol kgw -1 NaCl were studied at variable temperatures: 5-30°C for Mg 2+ and 10-50°C for Ca 2+ . Under appropriate conditions, the ternary complexes (MnUO2(CO3)3 (4-2n)-with n = 1 for Mg, n = {1;2} for Ca) were identified by time-resolved laserinduced luminescence spectrometry. Their pure spectra at 50°C for CanUO2(CO3)3 (4-2n)-and 30°C for MgUO2(CO3)3 2-were recovered by multivariate curve resolution alternating least squares analysis. Approximation models were tested to fit the experimental datathe equilibrium constants of complexation measured at different temperaturesand deduce the thermodynamic functions, i.e., enthalpy, entropy, and heat capacity. The weak influence of temperature on complexation constants induces large uncertainties on thermodynamic functions. Assuming the enthalpy is constant with temperature using the Van't Hoff equation, the first stepwise complexation of UO2(CO3)3 4-by Ca 2+ is estimated to be slightly endothermic with Δr VH H°m = (4.9 ± 5.9) kJ mol -1 , while the second stepwise complexation of CaUO2(CO3)3 2- by Ca 2+ with is slightly exothermic Δr VH H°m = -(6.2 ± 13.3) kJ mol -1 . In contrary to Ca 2+ , the complexation of UO2(CO3)3 4-by Mg 2+ is slightly exothermic with ∆r VH H°m = -(11.7 ± 10.3) kJ mol -1 . These values are not significantly different from nil inasmuch as the uncertainties are important due to a weak dependence of log10K° values. The entropic character of the complexation is verified as TΔrS°m = (36.5 ± 6.0) kJ mol -1 for the first stepwise complexation of UO2(CO3)3 4-by Ca 2+ , TΔrS°m = (13.8 ± 13.3) kJ mol -1 for the second stepwise complexation of CaUO2(CO3)3 2-by Ca 2+ , and TΔrS°m = (14.2 ± 10.3) kJ mol -1 for the complexation of UO2(CO3)3 4-by Mg 2+ . The energetics of complexation and sensitivity analysis of model solubility. 20 The formation constants of these complexes have been determined around standard temperature using different techniques, e.g., time-resolved laser-induced luminescence spectrometry (TRLFS), [21][22][23][24][25] Ca 2+ selective electrodes, 18 anion exchange, 26,27 and ultravioletvisible (UV/Vis) absorption spectroscopy. 28,29 Their molecular structures have also been well defined by EXAFS 19,21,22,30,31 and molecular simulations. 32,33 In recent works, [34][35][36] we have reported the determinations of the stepwise thermodynamic constants log10K°n for the equilibria,

where M 2+ represents the divalent alkaline earth free ion Mg 2+ and Ca 2+ , and the cumulative thermodynamic constants log10β°n.1.3 for the equilibria,

)-(n = 1 for Mg, and n = {1;2} for Ca)

in NaCl and NaClO4 electrolytes using TRLFS. We also obtained their specific interaction coefficients ε(MnUO2(CO3)3 (4-2n)-,Na + X -) (X -being Cl -and ClO4 -) using the specific ion interaction theory (SIT) that were absent from compilation of data [37][38][39] see our previous works [34][35][36] and references therein. Using these data, the thermodynamic functions of the ternary complexes under standard conditions are more complete and the speciation modelling should be more reliable.

However, very scarce data at varying temperature are available for MnUO2(CO3)3 (4-2n)- complexes. Endrizzi and Rao 18 measured the complexation of CanUO2(CO3)3 (4-2n)-complexes at 25°C using calorimetric titration. They concluded that the successive formations are entropy driven, and the contribution from enthalpy is rather small with ΔrH° = -(7.6 ± 5.8) and (0 ± 7) kJ mol -1 for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), respectively, and statistically no different from nil. Similar conclusions were reported using calorimetry by Jo, et al. 28 with ∆rH°m = -(1.2 ± 0.3) and -(3.1 ± 3.0) kJ mol -1 for CaUO2(CO3)3 2-for Ca2UO2(CO3)3(aq), respectively. The latter authors reported the equilibrium constants at variable temperatures (10-70 °C) using TRLFS, Ca 2+ selective electrode, UV/Vis absorption spectroscopy, and calorimetry. However, the experimental performance of the used techniques was limited to 70°C. For this reason, the experimental results measured in the temperature range of 10-50°C, though incomplete, were applied to the modified isoelectric reactions, as defined in Eq. V -15 and Eq. V -16, in order influence of temperature and ionic strength on MnUO2(CO3)3 (4-2n)-complexation will be analysed for particular scenarios in combination with our previous data. [34][35][36] V -3 Experimental Section

Sample preparation

The stock solution of U(VI) was prepared by dissolving high purity U3O8 in 37% hydrochloric acid (Sigma-Aldrich, ACS reagent). Inductively coupled plasma mass spectroscopy (ICP-MS) was used to measure the stock concentration of U(VI). The concentration of U(VI) in all studied solutions was fixed at 50 μmol kgw -1 . The ionic strength was maintained at 0.1 mol kgw -1 with anhydrous NaCl (Sigma-Aldrich, ACS reagent, ≥ 99%)

and Millipore deionized water (Alpha-Q, 18.2 MΩ cm). The addition of freshly prepared NaHCO3 (Analytical Grade, Fisher) solution was previously calculated considering the equilibrium between the atmospheric CO2 and the aqueous pH value at 25°C. Calcium chloride (CaCl2, Sigma-Aldrich, ACS reagent, ≥ 99%) and magnesium chloride (MgCl2, Sigma-Aldrich, ACS reagent, ≥ 99%) were used to control the concentrations of alkaline earth metal ions.

Freshly prepared NaOH (Analytical Grade, Fisher) and diluted HCl solutions were used for the pH adjustments. The samples were under saturated with respect to uranium (UO3:2H2O or β-UO2(OH)2), 39 and calcite 41 (for Ca-series) or magnesite 42 (for Mg-series). No obvious precipitation was observed throughout the experiments. The preparation of all samples was carried out by weighing under aerobic conditions.

The pH measurements were done using a combined hydrogen ion-sensitive glass electrode

Equilibrium constants and ionic medium corrections

The stoichiometry of MnUO2(CO3)3 (4-2n)-complexes is identified by the so-called slope analysis, which has been successfully applied in our previous studies, [34][35][36] to calculate the formation constants in NaCl and NaClO4 electrolytes of different ionic strengths. The pH values and Mg 2+ /Ca 2+ concentrations are changed simultaneously in spectroluminescent titration i) to avoid the occurrence of UO3:2H2O(cr) 38,39 and calcite (CaCO3) 41 or magnesite (MgCO3) 42 compared to the typical pH-fixed experimental methodology, and ii) to allow the formation of ternary species to a great magnitude in solution.

Successive formations of MnUO2(CO3)3 (4-2n)-are described with the conditional equilibrium constants Kn.1.3 expressed as follows,

where squared brackets indicate concentration in molality (mol kgw -1 ).

The logarithmic transformation of Eq. V -2 in Eq. V -3 allows calculating the log10Kn.1.3 value. The slope n in Eq. V -3 stands for the stoichiometry of Mg 2+ /Ca 2+ , Rf is the concentration ratio between complexed and non-complexed UO2(CO3)3 4-by Mg 2+ /Ca 2+ . The intercept log10Kn.1.3 refers to the conditional stepwise formation constants of MnUO2(CO3)3 (4-2n)-at a given ionic strength,

which can be converted to luminescence intensity supposing it is proportional to the U(VI) concentrations.

Here FI0(MnUO2(CO3)3 (4-2n)-) is the fluorescence intensity extrapolated at D = 0.

FI0(UO2(CO3)3 4-) is the one of the non-complexed UO2(CO3)3 4-at pH 9, and α represents the Ringböm coefficient 44 of UO2(CO3)3 4-. The variation of UO2(CO3)3 4-with pH at [M 2+ ] = 0 conditions was calculated to assess the value of αthe reader is referred to previous papers for detailed explanation and specific calculation. [34][35][36] The α values were calculated using the thermodynamic data (log10K° and ∆rH°m) from Table V -S3 of the SI, 38 and specific ion

V -5 Practical applications

Case of nuclear waste repositories

The Boom clay and the Callovo-Oxfordian (COx) clay, identified as high-level long-lived radioactive waste repositories in Belgium and France, 64,65 respectively, had been characterized

with the chemical compositions of equilibrium waters given in de Craen, et al. 9 and in Gaucher et al. 11 In our previous calculations using the experimentally determined constants log10K°n.1.3 and ε at near standard conditions, 34,36 it has been concluded that the high carbonate concentration in the Mol reference water in the Boom clay contributes to the formation of MnUO2(CO3)3 (4-2n)-at 25°C, though U(IV) should dominate the solubility and control the mobility of uranium because of the reducing redox potential (Eh = -274 mV, or P(O2)= 10 -68.67 atm). [66][67][68] In the case of the COx equilibrium waterthe solution A in Gaucher, et al. 11

V -8 Supporting Information

Chapter VI. Implications of Newly Derived Thermodynamic Data on the Predominance Diagrams of the Ca-U(VI)-CO 3 Systems at Room

Temperature.

VI -1 Context

In a previous study, Reiller and Descostes 1 proposed predominance rangesin a activity of SO4 2--pH domain at constant log10a(Ca 2+ ) log10a(Mg 2+ ) and P(CO2)of the (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes, based on speciation calculation of water compositions relevant of uranium mining activities and of other varying occurrences, using PhreeqC database output file from PRODATA 1.1.0.4 database. In the framework of PRODATA, the values of formation constants for CanUO2(CO3)3 (4-2n)-, SrUO2(CO3)3 2-and BanUO2(CO3)3 (4-2n)-were taken from Dong and Brooks 2 , and the one of MgUO2(CO3)3 2-was chosen from Dong and Brooks. 3 In chemical modelling, the evolution of log10β(MnUO2(CO3)3 (4-2n)-) as a function of ionic strength is important to be specified as the final speciation results can show significant differences due to different choices of activity corrections. However, the determination of specific ion interaction coefficients ε for these ternary complexesparticularly (Mg/Ca)nUO2(CO3)3 (4-2n)-complexesdid not receive sufficient attention and only the value of ε(MgUO2(CO3)3 2-, Na + ) was reported at the time of publication. 3 Therefore, ε(MgUO2(CO3)3 2-, Na + ) = -3.2 kgw mol -1based on Dong and Brooks 3 and ε(CanUO2(CO3)3 (4-2n)-, Na + ) = 0because of non-availabilitywere taken in the theoretical speciation calculations in Reiller and Descostes. 1 The main objective of this chapter is to examine the main species in the natural water compositions linked to radioactive waste management (RWM), uranium mining activity, and of other more general interest related to natural uranium with the newly determined thermodynamic data for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes obtained in previous chapters.

The selected water conditions particularly cover those reported for:

(a) waters in underground research laboratories (URL) for RWM, e.g., clay porewaters (Boom clay in Belgium, 4 Callovo-Oxfordian 5 and Tournemire 6 clays in France, Mont-Terri clay in Switzerland 7,8 ) and granitic groundwaters (Äspö site, the Forsmark and the Laxemar-Simpevarp areas in Sweden [9][10][11] );

Table VI -1 Occurrence of inorganic dissolved constituents in groundwater. 26 Major constituents (1.0 to 1000 mg L -1 )

Secondary constituents (0.01 to 10 mg L -1 )

Minor constituents (0.0001 to 0. For this exercise, a number of natural water compositions were collected from the literature available. The measurement of the geochemical parametersi.e. pH value, redox potential, salinity, alkalinity, temperature, etc.important characteristics for speciation modelling were taken for each collected sample. Generally, these natural waters are dominated by the following ions: Na + , Ca 2+ , Mg 2+ , K + , Cl -, NO3 -, HCO3 -, CO3 2-, and SO4 2-that may be attributed to the chemical weathering process and the leaching of minerals. Among these ions, chloride concentration is the most suitable parameter to describe water salinity because it is weakly lost from solution by precipitation or adsorption. The concentration of uranium in the collected samples highly depends on the sampling regions. In most natural waters, uranium concentration is between 4.2 10 -10 M and 4.2 10 -9 M. 27 In Europe, the average uranium concentration in continental surface water is estimated to be on the order of 0.3 μg L -1 (≈ 1.26 10 -9 M) while it is on the order of 0.5 μg L -1 (≈ 2.1 10 -9 M) worldwide. 28 In river waters, it can extend over four orders of magnitude with average content from 0.02 to 6 μg L -1 (≈ 8.40 10 -11 to 2.52 10 -8 M). 29,30 In France, the uranium concentration in river waters is about 0.44 μg L -1 (≈ 1.85 10 -9 M) in sedimentary areas, and expected to be a little lower than 0.15 µg L -1 in the Hercynian shelf and alpine areas. 31 However, in mines the uranium concentration in groundwater can range from 6.3 10 -8 to 1.68 10 -6 M. 32 Two water samples were collected from the Lodève site by ORANO Mining and their chemical compositions were analyzed. With this information, it is possible to perform a theoretical speciation calculation and then compared with the luminescence of the dominant species experimentally observed by TRLFS. Table VI -2 lists the chemical constituents for the investigated water samples: one was the "untreated" water initially at pH = 8.2 and the other one was initially acidified to pH 6.5 after sampling.

VI -3 Description of calculations of uranium speciation in the selected samples

The choices of the chosen software, the database file and activity correction would have a particular effect on the simulation results. The calculations in this chapter were performed using the PHREEQC version 3.4.0.14000 33,34 for the speciation calculation and Phreeplot 1.0 35,36 for the predominance diagrams. According to the SIT definition in PhreeqC, activity correction adopted in this work, the formation constants at infinite dilution log10β°, the thermodynamic functions ∆fHm° and SIT parameters, specific ion interaction coefficients ε for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes, were implemented in the input file, and the Thermochimie 9b database, 37 which shares the same body of thermodynamic data with PRODATA 1.1.0.4 database was used. the system Ca 2+ -U(IV/VI)-Na + -Cl --OH --HCO3 --H2O(l) as a function of partial pressure of oxygen. The activity diagrams log10(a(HCO3 -) a(H + )) vs. log10(a(Mg 2+ ) 0.5 a(Ca 2+ ) 0.5 /a(H + ) 2 )

VI -4 Description of activity predominance diagrams

have been chosen for the presentation due to the number of variables affecting this system and the non-linear relationship among them. It allows to represent the variation of H2CO3 and (Mg-Ca)Ca(OH)2. The calculations were conducted at fixed uranium concentration of 10 -8 mol kgw -1 at Im = 0.1 mol kgw -1 NaCl with P(O2) varying from 10 -73 to 10 -51 atm, which covers a wide range of actual underground research laboratory for the radioactive waste managemente.g., of P(O2) = 10 -73 atm for the Äspö site in Sweden up to, P(O2) = 10 -67 atm for the Callovo-(≈ 6.30 10 -8 M)the World Health Organization-proposed restrictive guideline 38 around the Cigar Lake mining site (Saskatchewan, Canada), 14 the Oklo uranium ore deposits, from several ORANO mining sites in an ISR context, 39,40 from the European project FOREGS, and sampled in the Semipalatinsk nuclear test site. 18 However, concentrations greater than 10 -6 M even up to the order of 10 -4 M can occur in several water compositions associated with uranium ore deposits, such as in the former uranium mine at Cluff lake (Canada), 21 in the U ore deposits within the Lodève Basin in South Central France, 6,13 and in Kazakhstan. Some of these groundwaters containing high uranium concentration are strongly reducing, and contain important sulphide but very low calcium concentrations. The speciation distribution for the selected aqueous solutions was carried out using PhreeqC. The activity of important ions -log10a(HCO3 -), log10a(H + ), log10a(Ca 2+ ), log10a(Na + )and the percentage of the major species were generated in the output files. Saturation index of calcite and schoepite was also calculated, which is key parameters in geochemical studies and indicates how far the solution is from thermodynamic equilibrium with respect to these minerals. In this study, we only consider the water compositions that are showing CanUO2(CO3)3 (4-2n)-as major complex. The selected water compositions are plotted in log10(a(HCO3 -) a(H + )) vs. log10(a(Ca 2+ )/a(H + ) 2 ) and log10(a(Na + )/a(H + )) vs. log10(a(Ca 2+ )/a(H + ) 2 ) activity predominance diagrams calculated at fixed uranium concentration of 100 μmol kgw -1 and P(O2) = 10 -0.6788 atm at 25°C considering the activity of water to be unity. The choice of axis ranges in Fig. VI -3 is based on the analysis of the frequency of the log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(HCO3 -) a(H + ))-log10(a(Na + )/a(H + ))

parameters. The most of selected samples within the range of log10(a(Ca 2+ )/a(H + ) 2 ) from 10 to 13, of log10(a(HCO3 -) a(H + )) from -11 to -9.5, and of log10(a(Na + )/a(H + )) from 4 to 6. The precipitation of calcite and dolomite is not requested in calculation because several water were determined from the distribution of sample points in log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -) a(H + )) diagram with the interval confidence of 68%, corresponding to 1σ. Using a higher confidence level, i.e. 95% or 98%, would generate wider ellipsoids, which tend to produce a larger intersection between the two ellipsoids of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), thus a less clear distinction of their concentration domain.

To allow representation, Fig. VI -9 illustrates the confidence ellipsoids for three species using 2D images in different view angles. One can note that the ellipsoid of Ca2UO2(CO3)3(aq) is slightly narrower and more elongated then that of CaUO2(CO3)3 2-. The overlap of the calculated ellipsoids of CaUO2(CO3)3 2-, and Ca2UO2(CO3)3(aq) can nearly extend to the half of their individual volumes. Furthermore, the ellipsoid of Ca2UO2(CO3)3(aq) is located in the regions of higher calcium and carbonate concentrations compared with that of CaUO2(CO3)3 2-, which is consistent with the evaluation of equilibrium constants. Regarding to the concentration region of UO2CO3(aq), its ellipsoid exists at low log10(a(Na + )/a(H + )) and low log10(a(Ca 2+ )/a(H + ) 2 ), underlining that the occurrence of major UO2CO3(aq) in aqueous phase is due to an insufficient calcium concentration in solution of low ionic strength. Nevertheless, it is noteworthy that the pH value is also an important parameter as the samples with UO2CO3(aq) as major species are all slightly acidic of pH value between 6 and 7, which is unfavorable to the formation of CaUO2(CO3)3 2-, and Ca2UO2(CO3)3(aq) especially at low calcium concentration.

VI -6 Comparison of modelling and experimental results for the water samples

collected from the Lodève site

Modelling

With the detailed chemical compositions for the "untreated" water and the acidified water sampled from the Lodève site (Table VI -2), their uranium speciation calculations were performed using log10β°, ∆fHm° and SIT parameters ε for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes implemented in the input file for PhreeqC and using Thermochimie 9b database. The percentages of major uranium carbonate species, the activity of important ions as well as the values of log10(a(Ca 2+ )/a(H + ) 2 ), log10(a(Na + )/a(H + )), and log10(a(HCO3 -) a(H + )) are listed in

VI -7 Interim conclusion

This chapter examines the speciation calculation of uranium in actual situations covering radioactive waste management (RWM), uranium mining activity, and other natural water samples with the previously determined thermodynamic data for (Mg/Ca)nUO2(CO3)3 (4-2n)- complexes in this study implemented in ThermoChimie 9b database. It further represents a comprehensive and the most updated thermodynamic dataset, including the formation constants at infinite dilution log10β°, the thermodynamic functions ∆fHm° and SIT parametersspecific ion interaction coefficients ε for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes -, which can be applied in geochemical calculations covering a wide range of conditions relevant to nuclear waste disposal. Compared with previous thermodynamic modelling executed with the old data in literature, data gaps can be identified in this study concerning the assessment of U(VI) chemistry under high ionic strength or high temperature as those expected in disposal environment and deserve further attention.

Based on the repartition of major uranium species obtained in the speciation calculations for the selected water compositions, three ellipsoidal predominance regions of CanUO2(CO3)3 (4-2n)- complexes were constructed from the sample point-clouds in log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -) a(H + )) three-dimension space. The determination of ellipsoids for (Can)UO2(CO3)3 (4-2n)-complexes allows a quick identification of the major (calcium) uranyl carbonate species for a given natural water sample once its chemical composition is available. Thermodynamic calculations and TRLFS analyses were performed for two water samples collected from the Lodève site. The calculated dominant species corroborates well the experimental observations in terms of important luminescence characteristics: band positions and decay time. The measured luminescent properties for (Mg/Ca)nUO2(CO3)3 (4-2n)-complexes under different conditionstype of electrolyte, ionic strength, and temperaturedetailed in preceding chapters and partly verified by the in-situ samples, surely enrich the reference catalog of Mg/Ca-UO2-CO3 luminescence spectral fingerprints.

Chapter VII. General Conclusions

VII -1 Determination of log 10 β° and ε for Ca n UO 2 (CO 3 ) 3 (4-2n)-complexes in NaCl

The detection of luminescence signal for uranyl carbonate complexes by TRLS was successful in the first investigation at room temperature, regardless of the quenching effect of carbonates. Hypsochromic-shifted peak positions compared with those of uranyl ions in acidic conditions provided evidence for the inner-sphere substitution of carbonate groups for water molecules. The complexation of calcium with uranyl carbonate moiety was then studied by titration of calcium at varying pH values where the Ringböm coefficients were introduced to describe the variation of UO2(CO3)3 4-relative to pH and CO2(g). The degree of complexation was estimated by the variation of fluorescence intensity as a function of the free Ca 2+ concentration. By these means, the stoichiometry of ternary species was identified by the slope analysis, which provided an approach to investigate separately their properties.

Due to the noticeable quenching effects of chloride ions, the upper limit of aqueous electrolyte concentration was imposed to 1 mol kgw -1 NaCl. The characteristic luminescent properties-spectrum and decay time-were recorded for all the prepared samples at each investigated ionic strength. It was observed that successive complexation of Ca 2+ with UO2(CO3)3 4-resulted in gradually enhanced fluorescence intensity, better-resolved spectrum, and slower luminescence decay-time. In addition, the luminescence spectra were strikingly similar for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq), except slightly narrowed peak widths for Ca2UO2(CO3)3(aq). The distinctive differences between CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq)

were from their decay times. It was assessed that in 0.1 mol kgw -1 NaCl medium, CaUO2 As a result, the binding processes of Mg 2+ and Ca 2+ with UO2(CO3)3 4-should not be estimated to the binary stage alone. Still, the ion-pairs formations between Mg 2+ /Ca 2+ and the medium anion were supposed to have comparable quantitative effects on the complexation of Mg 2+ /Ca 2+ with UO2(CO3)3 4-, resulting in the close values of interaction coefficients for MUO2(CO3)3 2-in a given electrolyte.

Then the applications of the determined data for MUO2(CO3)3 2-to several representative geochemical contexts stressed out the necessity to account for the MgUO2(CO3)3 2-complex in theoretical speciation calculation. Particularly when the ratio of magnesium over calcium approaches or overpasses the ratio of the formation constants.

VII -4 Effect of temperature on the formation of MUO 2 (CO 3 ) 3 2-

The complexation of triscarbonatouranyl(VI) with Ca 2+ and Mg 2+ were then studied by TRLFS at Im = 0.1 mol kgw -1 NaCl in the temperature range 10-50°C and 5-30°C, respectively.

At a specified temperature, the spectra showed an increase in the emission intensity with increasing alkaline earth metal concentrations, which indicated the successive complexation of Mg 2+ /Ca 2+ to UO2(CO3)3 4-and allowed the determination of their formation constants. As temperature increases, the global amplitude of intensity suffered a pronounced decrease and in particular, at high temperatures, i.e., 40, 45, and 50°C, the spectra showed slight shifts to higher wavelengths at the beginning of the complexation that allowed sorting out individual spectra of involved species using singular value decomposition (SVD) and supplemental calculation. Different approximation models were adjusted to the evolution of formation constants of MnUO2(CO3)3 (4-2n)-as a function of temperature. We concluded that the entropy term contributes primarily to the formation of MgUO2(CO3)3 2-and CaUO2(CO3)3 2-and the enthalpy contribution to these reactions may be both assumed to be almost nil due to the high uncertainties in fitting results. The second complexation of CaUO2(CO3)3 2-by Ca 2+ contributed to an exothermic process with a less important entropy value for molecular structural reasons. VII -5 Analysis of M n UO 2 (CO 3 ) 3 (4-2n)-in the natural water compositions

A number of natural water compositions were collected from the literature available on the subject of uranium mining and environmental monitoring activities. We only considered the water compositions that are showing CanUO2(CO3)3 (4-2n)-as major complex after speciation calculation using the PHREEQC with all the previously determined thermodynamic constants -log10β°, ∆fHm° and SIT parameters, ε for Mg/CanUO2(CO3)3 (4-2n)-complexes implemented in PRODATA 1.1.0.4 database. The ellipsoidal boundaries of UO2CO3(aq), CaUO2(CO3)3 2-, and Ca2UO2(CO3)3(aq) were determined from the distribution of sample points in log10(a(Ca 2+ )/a(H + ) 2 )-log10(a(Na + )/a(H + ))-log10(a(HCO3 -) a(H + )) diagram. The determination of ellipsoids for CanUO2(CO3)3 (4-2n)-complexes is a suitable tool for quickly identifying the dominant species for a given natural water sample once its chemical composition is available.

Two water samples were collected from the Lodève site by ORANO Mining. The two natural occurring water samples both delivered a measurable luminescence signal, which is characteristic for uranyl carbonate luminescence. With their detailed chemical compositions, the dominant species determined by theoretical speciation calculations corroborated well the measured luminescent properties by TRLFS.