Flávia Maia

List of Figures

Model approach considering thermodynamic data for surface reactions from Table 39 and using the thermodynamic data for Ca-U(VI)-CO3 aqueous complexes (see Table 14). Solid Table 1. Average and extreme values of the mineralogical composition (wt%) of the host layer in borehole EST205 Meure/Haute-Marne site [START_REF] Gaucher | ANDRA Underground Research Laboratory: Interpretation of the Mineralogical and Geochemical Data Acquired in the Callovian-Oxfordian Formation by Investigative Drilling[END_REF] (Bradbury andBaeyens, 1997, 2004) (Guillaumont et al., 2003) and surface complexation reactions and respective stability constants (Bradbury and Baeyens, 2009;Marques Fernandes et al., 2015) and enthalpy values (unknown: to be determined) used on the predictive model

List of Tables

Résumé Étendu

La présence éventuelle de radionucléides dans l'environnement, due à l'utilisation de la production d'énergie nucléaire et à la manipulation généralisée des radio-isotopes dans la recherche, l'industrie, la médecine et l'agriculture, soulève un certain nombre de préoccupations environnementales. Il est essentiel dans ce contexte de comprendre le comportement chimique des radionucléides pour permettre une prévision fiable de leur migration dans le milieu naturel.

En France, la formation argileuse du Callovo-Oxfordien (COx), située en Meuse / Haute-Marne (Est de la France), est à l'étude comme barrière géologique pour le projet de stockage des déchets radioactifs « Cigéo : Centre industriel de stockage géologique ». Cette formation argileuse, de plus de 130 m d'épaisseur et à 500 m de profondeur, est caractérisée par une faible perméabilité ce qui limite la circulation de l'eau, principale cause de la dissémination des radioéléments. De plus, il s'agit d'une zone qui présente une bonne homogénéité sur plusieurs km 2 (minéralogie, propriétés physicochimiques) [START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF]. Enfin, elle est caractérisée par un assemblage minéralogique riche en phases argileuses qui jouent un rôle important dans le contrôle des propriétés physico-chimiques et le ralentissement de la migration des radionucléides compte tenu de leurs importantes capacités de rétention.

L'environnement géologique au voisinage d'un stockage de déchets nucléaires sera évidemment soumis à des perturbations (par exemple thermiques, hydrauliques, mécaniques et chimiques, connues sous le nom de THMC) et à différents niveaux de radioactivité au cours du temps. L'une des perturbations primordiales durant les premières années est l'augmentation de la température autour du stockage, favorisée par les désintégrations radioactives au sein des déchets de haute activité à vie longue. Dans la roche argileuse hôte, cet effet peut non seulement modifier la composition chimique de l'eau interstitielle (pH, potentiel d'oxydoréduction et pression du gaz), mais aussi provoquer des processus d'interaction eau-roche-gaz (par exemple : dissolution, précipitation), voire modifier les propriétés de sorption.

Actuellement, dans les installations françaises, le combustible usé est composé en plus de la matière fissile (oxydes d'uranium et oxydes mixtes d'uranium et de plutonium) des produits de fission et des actinides mineurs responsables de plus de 98 % de la radioactivité. Ces derniers, considérés comme des déchets, sont séparés pour être conditionnés au sein de matrices de verre par des procédés de vitrification [START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF]. Le combustible usé, qui n'est pas retraité, n'est pas considéré comme un déchet. Le degré de radioactivité de ces combustibles usés est comparable à celui des déchets vitrifiés. L'interaction du rayonnement α émis par les radio-isotopes avec l'eau peut conduire à des réactions provoquant la corrosion et la dissolution du verre, libérant ainsi les radioéléments (notamment U, Pu, etc…). La mobilité de l'uranium dans la géosphère est intrinsèquement liée à sa solubilité qui dépend de son degré d'oxydation (communément en conditions naturelles U (IV) ou U (VI) ) et à la stabilité des espèces aqueuses présentes dans l'eau porale de la formation argileuse (Runde, 2000).

Les résultats d'exercices de modélisation géochimique, qui s'appuient sur les données thermodynamiques issues du projet international coordonné par l'Agence de l'Energie Nucléaire (NEA-TDB ; Guillaumont et al., 2003), prédisent l'existence de l'uranium au degré d'oxydation U (IV) dans les conditions physico-chimiques représentatives de l'eau porale des argilites à l'état naturel. Au degré d'oxydation U (IV) , l'uranium est peu soluble et les espèces dissoutes associées peuvent réagir fortement avec les surfaces argileuses des argilites du Callovo-Oxfordien ce qui rend cet élément peu mobile. Cependant, des données publiées ont clairement établi l'existence de complexes ternaires Ca-U(VI)-CO3 (Bernhard et al., 1996(Bernhard et al., , 2001 ; ;Kalmykov et Choppin, 2000 ;Dong et Brooks, 2006 ;Lee et Yun, 2013 ;Endrizzi et Rao., 2014 ;Jo et Yun, 2017) dans les conditions physico-chimiques des argilites (pH ≈ 7,2 et Eh ≈ -200 mV). Ainsi, la présence de ces complexes ternaires en solution peut conduire à stabiliser l'uranium au degré d'oxydation U (VI) en solution et donc augmenter la solubilité de l'uranium et une mobilité plus élevée du fait d'une rétention moindre sur les surfaces argileuses (Joseph et al., 2013 ;Vercouter et al., 2015 ;Tullborg et al., 2017).

Le travail de thèse s'inscrit dans ce contexte et a pour but de mieux comprendre le comportement de l'uranium dans la formation des argilites du Callovo-Oxfordien en prenant en compte de façon exhaustive toutes les données publiées notamment sur la stabilité des complexes ternaires d'uranyl. L'effet de la température (20 à 80 °C) sur ce comportement est particulièrement étudié.

La première partie du travail a pour but d'identifier et de quantifier l'effet de la température sur la formation des complexes ternaires Ca-U(VI)-CO3. Les processus de rétention de l'uranium sont ensuite étudiés à 20 °C, en se concentrant sur la sorption de U(VI) sur des minéraux argileux de référence (illite et montmorillonite), sur la fraction argileuse des argilites du Callovo-Oxfordien (illite et interstratifié illite/smectite) et au contact d'échantillon brut d'argilites du Callovo-Oxfordien. Le modèle de sorption développé à 20 °C, combiné avec les paramètres thermodynamiques qui caractérisent les complexes ternaires, a servi de base pour l'étude de l'effet de la température sur la sorption de U(VI) pour les divers systèmes argileux.

Effet de la température sur la stabilité des complexes ternaires dans le système « Ca-U(VI)-CO3 »

La mesure des constantes de complexation aqueuse à différentes températures est basée sur une méthode de compétition proposée par Dong et Brooks (2006). La méthode a été adaptée pour permettre l'étude en température (via l'utilisation d'une résine plus stable) et la concentration en uranium a été ajustée pour éviter tout problème de précipitation. Des tests préliminaires ont été réalisés pour garantir le bon fonctionnement du montage expérimental et pour évaluer l'équilibre entre la pression de CO2(g) et les solutions aqueuses.

Les résultats indiquent que les deux complexes aqueux ternaires étudiés ont des comportements différents en température. Avec l'augmentation de la température, la formation du complexe anionique (CaUO2(CO3)3 2-) est défavorisée, alors que celle du complexe neutre (Ca2UO2(CO3)3(aq)) n'est pas influencée.

Ainsi, à partir des données expérimentales, les constantes de stabilité pour CaUO2(CO3)3 (Bradbury et Baeyens, 1997[START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF], 2011 ;Chen et al., 2014 ;Marques Fernandes et al., 2015). Le modèle 2SPNE SC/CE proposé par Bradbury et Baeyens (1997) a été utilisé pour décrire les interactions entre U(VI) et les phases argileuses modèles.

Si le système U(VI)/montmorillonite est bien décrit et paramétré (Bradbury et Baeyens, 1999 ;Marques Fernandes et al., 2012 ;[START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF]Tournassat et al., , 2018)), ce n'est pas le cas pour le système U(VI)/illite pour lequel l'influence des carbonates sur la rétention a été très peu étudiée. C'est pourquoi, un travail expérimental a été réalisé avec l'illite, en complément des études réalisées avec la fraction extraite du COx et des argilites du COx.

Le modèle proposé pour l'illite peut relativement bien expliquer les données expérimentales obtenues pour une large gamme de conditions expérimentales (pH = 6 à 7,3 ; pCO2 de 10 -3,5 à 10 -2,1 ; NaNO3 0,1M et composition chimique de l'eau porale synthétique des argilites du Callovo-Oxfordien). Il considère les paramètres de la montmorillonite pour les espèces hydrolysées de surface. Cependant, une nouvelle constante pour le complexe de surface ≡SsOUO2(CO3)3 2-est proposée. La valeur de la constante de sorption apparaît plus élevée pour l'illite (log10 K = 17,5) que celle proposée pour la montmorillonite (log10 K = 15,5, Marques Fernandes et al, 2012).

Le modèle permet également de décrire de manière convenable le comportement de U(VI) en présence de la fraction argileuse des argilites du Callovo-Oxfordien (illite et I/S), bien qu'un léger décalage ait pu être observé entre l'expérience et le calcul en absence de carbonates. Dans la mesure où le modèle proposé fonctionne correctement dans les conditions physico-chimiques représentatives de la formation du Callovo-Oxfordien (présence de carbonates), aucun changement de paramètre n'a été effectué et l'hypothèse de décrire I/S comme deux pôles d'illite et de montmorillonite est considérée comme valable.

Enfin, le modèle proposé a été appliqué avec succès pour décrire les données de sorption obtenues pour le système U(VI)/argilites du Callovo-Oxfordien. Il permet également de décrire le comportement de U(VI) en présence d'autres roches argileuses, telles que les argiles de l'Opalinus (Amayri et al., 2016 ;Joseph et al., 2013 ;Marques Fernandes et al., 2015)

Chapter 1 Introduction

The possible existence of radionuclides (RN) in nature, due to the increasing use of nuclear power generation and the widespread use of radioisotopes for beneficial purposes in research, industry, medicine and agriculture, raises a number of environmental concerns. Therefore, it is essential to understand the chemical behaviour of these radioelements for reliable prediction of their migration.

The disposal of radioactive waste has been discussed for many decades because of the special status that it has with the public, leading to implementation and development programmes that are typically mobilised by research and development efforts, resulting in a number of alternative disposal options.

The options reflect the specifics of national legislations, geological differences and variations in the amount and characteristics of different waste types (IAEA, 2016).

Countries such as Argentina, Belgium, Canada, Czech Republic, Finland, France, Japan, the Netherlands, Republic of Korea, Russia, Spain, Sweden, Switzerland, the UK, and the USA have chosen to dispose of radioactive waste in deep geological underground repositories. The usage of these repositories is generally reserved for intermediate-and high-level waste (ILW and HLW) and for long halflife radionuclides. The waste is mostly generated by the nuclear fuel cycle, mainly in the form of spent fuel and a variety of other waste types, as a consequence of the reprocessing cycle (such as vitrified waste and bituminous wastes). The rationale behind deep geological disposals is based on storing the waste packages in specifically engineered facilities located deep underground (often 400m to 600 m deep) in geological contexts. The main types of host-rock formations used are sedimentary (e.g. claystones), crystalline (e.g. granite) and salt domes [START_REF] Altmann | Geo"chemical Research: A Key Building Block for Nuclear Waste Disposal Safety Cases[END_REF].

Many countries have chosen to use claystone formations (e.g. Opalinus Clay in Switzerland; Boda Claystone Formation in Hungary; Boom Clay in Belgium; Callovo-Oxfordian Clay Formation in France) as deep geological repositories and these are the ones we are interested in the present work These countries have dedicated large research efforts on choosing an optimal type of rock (e.g. through the analysis of mineralogy and physical-chemical and mechanical properties) and a favourable geological context (e.g. depth, tectonic stability and groundwater systems). Moreover, research has also characterised the properties of the host-rock which determine its capacity to limit radionuclides transfer to the geosphere (such as hydraulic permeability, solute diffusion characteristics and retention processes).

The deep repository allows the confinement of the radioactive waste in a facility that ensures longterm safety through a multiple-barrier system to isolate the waste from the hydrosphere and biosphere.

As a result, future generations are protected from radioactive risk and toxicity. The multi-barrier system includes from the single-waste container to the near-and far-field: spent fuel disposal cells; concrete/metal canisters; backfill materials composed by cements and engineered swelling clays as bentonite; and natural host-rock as geological barrier (see example in Figure 1, showing the Andra disposal concept in claystone host-rock). The safety of radioactive disposal results from the performance assessment simulations that are based on modelling the processes that govern the radionuclide speciation and migration in natural systems [START_REF] Mckinley | Deep Geological Disposal of Radioactive Waste[END_REF]. The quantitative description of the radionuclides migration relies on a detailed knowledge of the system's geochemistry across the potential migration pathways (e.g. aqueous speciation, sorption, incorporation, precipitation) (Figure 2). Here, two main parameters have to be considered: the mobility of the radionuclides in solution and the retention that describes the affinity of radionuclides to the mineral surface, which also help in limiting their mobility.

The geological environment in a nuclear waste repository will also be subjected to disturbances (e.g. thermal, hydraulic, mechanical and chemical, known as THMC) and different levels of radioactivity over time. One of the consequences of this is the increase in temperature of all repository promoted by the radioactivity of the long half-life HLW. In the host-rock, this effect may not only change the porewater chemical composition (aqueous speciation, pH, redox potential and gas pressure) but also the water-rock interaction processes (dissolution, precipitation and sorption). Flávia Maia | IMT Atlantique Nantes

For this reason, it is essential to understand the processes of retention, through the analysis of sorption, and determine the stability of the surface complexes over a wide range of temperatures in order to predict the mobility of the radionuclides (Halter, 1999).

Context

In France, the Callovo-Oxfordian claystone formation (so-called COx) is under study as a geological barrier in the "Cigéo" (Centre Industriel de Stockage Géologique) radioactive waste disposal project, located in Meuse/Haute-Marne site (Eastern France). This argillaceous formation is characterised by low permeability, mineralogical homogeneity and low variation of other physicochemical properties in a large extension [START_REF]Dossier 2005: Andra Research on the Geological Disposal of High-Level Long-Lived Radioactive Waste. 4p[END_REF].

The Callovo-Oxfordian claystone, together with an engineered barrier system (EBS), comprises a mineralogical assemblage that plays an important role on the physical-chemical control and retardation of radionuclides migration to the natural environment.

At present, in French facilities the spent fuel is composed by uranium oxides and mixed uraniumplutonium oxides, which are conditioned in glass matrices by vitrification processes. The diffusion of α-radiation promoted by radioisotopes and water causes corrosion and dissolution of the spent fuel releasing radionuclides.

Uranium is one of the main components of the radioactive waste and it is usually found in its fourth oxidation state mainly in solid form in the repository's surrounding. Its mobility in the environment is intrinsically related to its solubility, which is limited by a) the stability of the solubility-controlling mineral, and b) the stability of the aqueous species (Runde, 2000).

In the surrounding environments of the repository, in contact with the groundwater, U(IV) may oxidise to U(VI), which results in stable and mobile complexes. The U(VI) may be present even under reducing conditions as a result of the formation of the ternary calcium-uranyl-carbonate aqueous complexes (Ca-U(VI)-CO3). As shown in Figure 3, the uranium released from vitrified waste in uranyl form will be controlled by the retention properties of the host-rock minerals (clay fraction) and by the chemical composition of the claystone porewater. In this study, we will focus on the sorption processes and the aqueous complexation of U(VI) in solution. The processes described in Figure 3 are important for the performance assessment of nuclear waste disposal under clay rock conditions. According to the scenarios developed for the Cigéo project, the temperature of the Callovo-Oxfordian formation must reach 90 ºC in the short-term and not more than 70 ºC in the long-term (for at least 1000 years) [START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF]. Changes in temperature can potentially induce changes in uranium behaviour (e.g. solubility, retention and mobility) through physical-chemical changes of the host-clay rock and porewater composition.

Thus, it becomes relevant to evaluate the effect of temperature on the chemical behaviour of U(VI).

Characterisation of the Callovo-Oxfordian Formation System

The Callovo-Oxfordian claystone is a geological formation belonging to the Paris Basin (Figure 4). The Paris basin is composed of alternating claystone and limestone layers in a simple and regular geometric structure, deposited during a stable marine environment in the Jurassic period. The Callovo-Oxfordian unit is surrounded by massive carbonate units (Dogger and Oxfordian limestones) and is a predominantly clay-rich layer, 130m to 160m thick and located between 500m and 630m deep in the transition zone. The unit is divided in three sedimentary sequences: the central part is the most clayey zone (up to 60% clay minerals); whereas the top and bottom are characterised by a carbonate enrichment, related with the transition to the limestone units. At the transition zone and beyond, the unit shows very little lateral variation in lithology.

In terms of mineralogy, the Callovo-Oxfordian claystone formation is composed of phyllosilicates (20-60 wt%, mainly illite and interstratified illite/smectite, kaolinite, micas and chlorite), tectosilicates (10-40 wt%, mainly quartz and feldspars), carbonates (15-75wt%, mainly calcite and dolomite) and sulphides (0-3wt%, mainly pyrite). An example of a mineralogical composition of a borehole crossing Meuse/Haute-Marne site is provided in Table 1. The chemical composition of the claystone porewater is essential to assess the chemical and diffusive behaviour of the radionuclides, as well as the stability of the materials present in the nuclear waste disposal. Knowing the main chemical conditions of the clay-host rock porewater is thus an important step in predicting how the clay-host rock barrier will behave over time [START_REF] Altmann | Geo"chemical Research: A Key Building Block for Nuclear Waste Disposal Safety Cases[END_REF].

The chemical composition of the Callovo-Oxfordian porewater has been determined by in situ measurements and by complementary geochemical calculations taking into account the mineralogical composition. The latter is done due to the difficulty of sampling directly the porewater of COx claystone (or any other type of clay rock), where only small volumes can usually be obtained to allow a quantitative representation of the in situ conditions. However, improvements in characterisation techniques have made it feasible and, as a result, representative porewater samples have been obtained, as shown by [START_REF] Vinsot | In Situ Characterization of the Callovo-Oxfordian Pore Water Composition[END_REF]. The comparison of the chemical composition measured in situ and that calculated validates the geochemical model proposed by [START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF], taking into account the mineral's solubility and cationic exchange reactions (using ThermoChimie database, Giffaut et al., 2014 and[START_REF] Grivé | Thermodynamic Data Selection Applied to Radionuclides and Chemotoxic Elements: An Overview of the ThermoChimie-TDB[END_REF].

As previously mentioned, the temperature increase originated from the radioactivity of the HLW will change the porewater chemistry. This temperature effect was experimentally studied by using different mineralogy assemblages [START_REF] Beaucaire | A Thermodynamic Model for the Prediction of Pore Water Composition of Clayey Rock at 25 and 80°C -Comparison with Results from Hydrothermal Alteration Experiments[END_REF]Tournassat et al., 2008). Two systems were defined: a) complete mineralogical composition, representative of COx claystone, composed of quartz, illite, chlorite, carbonates (calcite, dolomite and siderite), pyrite and celestine; and b) reduced mineralogical composition characterised only by carbonates and sulphates (calcite, dolomite, siderite and celestine). The results of the chemical composition of the different in situ porewaters (25ºC) and at different temperatures considering both mineralogy assemblages are provided in Table 2 andTable 3 (Tournassat et al., 2008).

Scopes of the thesis work

In the past, geochemical modelling exercises applied to nuclear waste disposal usually predicted U(IV) as the most important redox state when using thermodynamic data sets, such as NEA-TDB (Guillaumont et al., 2003). Recent published data have clearly established the existence of Ca-U(VI)-CO3 ternary complexes (Bernhard et al., 1996(Bernhard et al., , 2001;;Kalmykov and Choppin, 2000;Dong and Brooks, 2006;Lee and Yun, 2013;Endrizzi and Rao, 2014;[START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF]. The presence of these complexes leads to an increase of the predicted uranium mobility (greater solubility and lower retention) (Joseph et al., 2013;Vercouter et al., 2015;Tullborg et al., 2017). For this reason, a special care is taken for uranium, since it is considered to be the most mobile actinide in the performance assessment.

Therefore, the aim of the present work is to better understand the behaviour of U(VI) in Callovo-Oxfordian formation, including how it is affected by temperature changes.

This work is divided into three main parts. In Chapter 2 we aim to identify and understand the dominating uranyl aqueous species at different temperatures, with a special focus on the thermodynamic properties of Ca-U(VI)-CO3 ternary complexes. In Chapter 3 we aim to understand the retention processes of uranium at 20°C, focusing on the U(VI) sorption in the most relevant clay minerals (illite and interstratified illite/smectite) of the Callovo-Oxfordian claystone system. In the last Chapter (Chapter 4), we used the model proposed at 20 ºC for the retention mechanisms of U(VI) to further study the effect of temperature on U(VI) sorption on Callovo-Oxfordian claystone systems at different temperatures.

Chapter 2 Speciation of Ca-U(VI)-CO 3 ternary complexes

Introduction

The evaluation of the migration of actinide elements from a deep underground repository requires reliable thermodynamic stability data on complex formation of actinide ions. The existence of ternary calcium-uranyl-carbonate aqueous complexes (Ca-U(VI)-CO3), and their impact on aqueous systems, has been studied extensively (Bernhard et al., 1996(Bernhard et al., , 2001;;Steward et al., 1996;Kalmykov and Choppin, 2000;[START_REF] Zheng | Influence of Calcium Carbonate on U (VI) Sorption to Soils[END_REF]Dong and Brooks, 2006;[START_REF] Kelly | X-Ray Absorption Spectroscopy Identifies Calcium-Uranyl-Carbonate Complexes at Environmental Concentrations[END_REF][START_REF] Meleshyn | Influence of (Calcium-)Uranyl-Carbonate Complexation on U(VI) Sorption on Ca-and Na-Bentonites[END_REF][START_REF] Prat | Uranium Speciation in Drinking Water from Drilled Wells in Southern Finland and Its Potential Links to Health Effects[END_REF]Joseph et al., 2011[START_REF] Joseph | Do Elevated Temperatures and Organic Matter Influence the U(VI) Diffusion through Argillaceous Rock? in Clay in natural engineered barriers for radioactive waste confinement[END_REF]Joseph et al., , 2013;;Lee and Yun, 2013;Endrizzi and Rao, 2014;Vercouter et al., 2015;[START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF]Tullborg et al., 2017).

Recently, predictive modelling performed by Vercourter et al. (2015) indicated that alkaline earth ternary carbonate species of M(II)-U(VI)-CO3 (where M(II) = Mg or Ca) dominate the uranyl speciation in natural water and clay conditions. Geochemical calculations carried out by Joseph (2013) and Tullborg et al. (2017) also indicated that the formation of calcium uranyl carbonates can stabilise uranium in solution and increase its mobility. Moreover, among the different actinides present in the radioactive waste, uranium is the most mobile. Therefore, the study of these complexes, and of their behaviour at temperatures higher than 25 ºC, is highly relevant for the knowledge of uranium behaviour in nuclear waste disposal. In this context, a brief overview of literature data on the system Ca/Mg-U(VI)-CO3 is provided below.

Ca-U(VI)-CO 3 system

The Ca-U(VI)-CO3 complexes were primarily identified by Bernhard and co-workers (1996) in seepage waters with high content of calcium and carbonates. TRLFS (Time-resolved laser induced fluorescence) spectroscopy analysis was used to identify these complexes. The authors detected that in the presence of calcium the fluorescence intensity of uranyl carbonates increases, while the uranyl carbonate species itself exhibits poor fluorescence, suggesting the formation of calcium-uranyl-carbonate aqueous species under such chemical conditions. The identification of these species and the uncertainty of the calculations have led to the on-going investigation of a better understanding of Ca-U(VI)-CO3 species formation. Consequently, different studies were developed to identify and validate the presence of Ca-U(VI)-CO3 aqueous species in different experimental conditions (Kalmykov and Choppin, 2000;Bernhard et al., 2001;Dong and Brooks, 2006;Lee and Yun, 2013;Endrizzi and Rao, 2014;[START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF].

The stability constants for Ca2UO2(CO3)3(aq) and CaUO2(CO3)3 2-can be described by Reaction 1and Reaction 2 and a summary of the literature that includes the data is provided in Table 4. In the NEA (Nuclear Energy Agency) chemical thermodynamics update for uranium (Guillaumont et al., 2003), a review on the available publications up to 2003 for those species was made. The authors analysed the studies from Bernhard et al. (2001) and Kalmykov and Choppin (2000). Guillaumont et al. (2003) pointed out some shortcomings in those experimental works, especially in relation to the ionic strength control of the system. For this reason, Guillaumont et al. (2003) did not support the selected stability constants for the formation of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq). However, the NEA review stated that the constant proposed in Kalmykov and Chopin (2000) and Bernhard et al. (2001) may be used as a guidance. For further information, a detailed discussion of these studies can be found in Appendix A in Guillaumont et al. (2003).

2Ca 2+ + UO 2 2+ +3CO 3 2-⇌ Ca 2 UO 2 (CO 3 ) 3 (aq) Reaction 1 Ca 2+ + UO 2 2+ +3CO 3 2-⇌CaUO 2 (CO 3 ) 3 2- Reaction 2
Later in 2006, Dong and Brooks determined the formation constants from anionic-exchange experiments with a calcium concentration range of 0 to 5 mM at pH = 8.1 and 0.1 M ionic strength by using NaNO3 as background electrolyte under room conditions. The authors pointed out that the constants agree with the expected trend: the binding constant of Ca 2+ to UO2(CO3)3 4-(Reaction 3) is larger than Ca 2+ binding constant to CaUO2(CO3)3 2-(Reaction 4).

Ca 2+ + UO 2 (CO 3 ) 3 4-i ⇌ CaUO 2 (CO 3 ) 3 2-
log10 Kº = 5.34 Reaction 3

i Note: UO 2 2+ +3CO 3 2-⇌ UO 2 (CO 3 ) 3 4-, Logβº013 = 21.84 according to data selection in Guillaumont et al. (2013).

Ca 2+ + CaUO 2 (CO 3 ) 3 2-⇌Ca 2 UO 2 (CO 3 ) 3 (aq)

log10 Kº = 3.52 Reaction 4

In 2013, Lee and Yun studied the chemical behaviour of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) species under atmospheric CO2(g) in weakly alkaline conditions using TRLFS measurements. Endrizzi and Rao (2014) studied the Ca-U(VI)-CO3 system by ion selective electrode potentiometry in 0.1 M of NaCl at pH 10 and [START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF] are studying spectroscopic properties and formation constants of the ternary Ca-U(VI)-CO3 complexes using TRLFS measurements. Although all these authors used different experimental approaches, the stability constants suggested for the formation of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) are in agreement with the ones reported by Dong and Brooks (2006) (see Table 4).

Besides the investigation of their thermodynamic properties, some studies have analysed the structure of Ca-U(VI)-CO3 complexes using EXAFS (Extended X-Ray Absorption Fine Structure) (Bernhard et al. 2001;[START_REF] Kelly | X-Ray Absorption Spectroscopy Identifies Calcium-Uranyl-Carbonate Complexes at Environmental Concentrations[END_REF] or molecular dynamic simulations [START_REF] Kerisit | Molecular Simulation of the Diffusion of Uranyl Carbonate Species in Aqueous Solution[END_REF].

Finally, some works used the stability constants from Dong and Brooks (2006)

Mg-U(VI)-CO 3 system

Besides the stability of Ca-U(VI)-CO3 complexes, only a few studies have been published concerning the stability of Mg-U(VI)-CO3 aqueous complexes [START_REF] Dong | Formation of Aqueous MgUO2(CO3)3 2-Complex and Uranium Anion Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium Exchange Mechanism onto an Exchange Resin[END_REF][START_REF] Geipel | Mixed Complexes of Alkaline Earth Uranyl Carbonates: A Laser-Induced Time-Resolved Fluorescence Spectroscopic Study[END_REF]Endrizzi and Rao, 2014;[START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF]. The stability of MgUO2(CO3)3 2-has been experimentally studied using different analytical techniques (see a summary in Table 5); the resulting values seem to be in reasonable agreement. On the other hand, the stability of Mg2UO2(CO3)3(aq), to the best of our knowledge, has only been reported recently by [START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF]. The constants provided by [START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF], as also reported by Dong and Brooks (2006) for the Cacounterpart, agree with the expected trend: the binding constant of Mg 2+ to UO2(CO3)3 4-(Reaction 5) is larger than the Mg 2+ binding constant to MgUO2(CO3)3 2-(Reaction 6).

Mg 2+ + UO 2 (CO 3 ) 3 4-⇌ MgUO 2 (CO 3 ) 3 2- log10 Kº = 3.96 Reaction 5 Mg 2+ + MgUO 2 (CO 3 ) 3 2-⇌ Mg 2 UO 2 (CO 3 ) 3 (aq) log10 Kº = 1.30 Reaction 6
Besides the investigation of their thermodynamic properties, [START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF] have also investigated the structure of Mg-U(VI)-CO3 complexes using EXAFS.

According to Endrizzi and Rao (2014) and [START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF], the main difficulty in determining the Mg2UO2(CO3)3(aq) stability constants is caused by its lower stability compared with the corresponding stability constant of Ca2UO2(CO3)3(aq).

Effect of temperature

Although the formation behaviour of the calcium-uranyl-carbonate ternary complexes is widely studied at room temperature conditions, there is a lack of thermodynamic data in conditions of increased temperature. The effect of temperature on Ca-U(VI)-CO3 complexes is not clear. Only one study (Endrizzi and Rao 2014) is published and Jo and Yun (2017) have shown preliminary results using a different experimental approach.

In 2014, Endrizzi and Rao determined the enthalpy of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) using microcalorimetry at ionic strength (I) 0.1 M of NaCl. The experimental data were fitted using HypDeltaH software; the enthalpy data were not corrected to infinite dilution (I = 0) since the effect of the media in the ionic strength ranged between 0 and 0.1 M, which was considered to be negligible. The reported enthalpy data are shown in Reaction 7 and Reaction 8.

Ca 2+ + UO 2 (CO 3 ) 3 4-i ⇌ CaUO 2 (CO 3 ) 3 2- ΔrHº = -8 ± 6 kJ/mol Reaction 7 Ca 2+ + CaUO 2 (CO 3 ) 3 2-⇌ Ca 2 UO 2 (CO 3 ) 3 (aq) ΔrHº = 0 ± 7 kJ/mol Reaction 8
The authors used an initial concentration of uranium of 20 mM at pH 10 -10.3 in NaCl 0.1 M. The high uranium concentrations are probably required for the correct performance of the microcalorimeter, however, it may lead to precipitation of U(VI) solids. Although the presence of solids is not reported in the original article, PhreeqC calculations performed in the context of the present work indicate that sodium uranate (as Clarkeite) may be oversaturated under those conditions (see in Appendix 4 Figure D. 4).

Jo and Yun (on-going) have been working on the spectroscopic properties and formation constants of the ternary Ca-U(VI)-CO3 complexes at different temperatures. Their work consists in determining the complexes' thermodynamic data, including stability constants and enthalpy values, using TRLFS measurements at a range of temperatures between 10 and 80ºC. Recently, [START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF] showed preliminary data relative to the stability constants (at I = 0.1 M) for Reaction 1 and Reaction 2 at different temperatures. Their results are shown in Table 6. In the calculation of the stability constants there is a clear temperature effect, where increasing the temperature leads to an increase of the stability constants; however, these results should be taken with caution as the work is still on-going and exhaustive experimental data are not available yet.

Based on TRLFS analysis, the authors suggested that uranyl hydrolysis was observed at low concentrations of calcium, where CaUO2(CO3)3 2-is predominant. For this reason, the authors only suggested enthalpy values for the Ca2UO2(CO3)3(aq) species [START_REF] Jo | Spectroscopic Properties and Formation Constants of Ternary Ca-UO 2 -CO 3 Complexes at High Temperatures[END_REF]. The stability constant corrected for infinite dilution using SIT was determined as an average value of 30.64 ± 0.17 for Reaction 2, and the enthalpy of reaction was -35 ± 18 kJ/mol.

i Note: UO 2 2+ +3CO 3 2-⇌ UO 2 (CO 3 ) 3 4-, ΔrHº = -39.20 according to data selection in Guillaumont et al. (2013).

Following the discussion above, and the relevance of the Ca-U(VI)-CO3 complexes for nuclear waste disposal in clay conditions, it seems necessary to further study the effect of temperature in the formation of these species.

Determination of the thermodynamic data of Ca-U(VI)-CO 3 aqueous species

From the discussion in the previous section, the existence and stability of the Ca-U(VI)-CO3 aqueous species at 25 ºC is confirmed. Nevertheless, the effect of temperature on the formation of these complexes is not yet clear.

Stability constants are mostly determined by equilibrium methods, where it is assumed that the activity coefficients of all species is effectively constant by the use of a suitable ionic medium, so that the law of mass action can be applied to the concentrations of reacting species (Rossotti and Rossotti, 1961).

Most aqueous complexes form reaching the equilibrium rapidly and their stability constants can be determined by direct or indirect methods. In direct methods, such as TRLFS or potentiometry, the aqueous species of interest is directly measured. For the case of the indirect methods, the stability constants are determined based on the analysis of the variation of element concentration on a biphasic system (such as solubility, ionic exchange and partition experiments) as a function of different parameters (e.g. temperature, solid-liquid ratio) that change the chemical system.

The objectives of the present work are:

1) To design an experimental method for the ionic exchange experiments performed in a different range of temperatures, where the control of temperature and partial pressure of CO2 of the system are our main concern; 2) To evaluate the reliability of the stability of Ca-U(VI)-CO3 aqueous species at 25ºC by corroborating it with the literature data and to provide confidence to the experimental methodology and modelling approach. 3) To study the effect of temperature between 20 and 80ºC on the formation and stability of Ca-U(VI)-CO3 complexes.

Ionic Exchange Method

In the present work, the ionic exchange method was used to determine the stability constants of the Ca-U(VI)-CO3 complexes and the effect of temperature on their stability. This indirect method was previously used by Dong and Brooks (2006) to determine the stability constants of the Ca-U(VI)-CO3 system at 25ºC under atmospheric conditions. The approach is based on the distribution of a metal (uranium in this case) between an ionic-exchange resin and the aqueous phase (Reaction 9).

M aq ⇌ M resin Reaction 9

In this reaction, M is the solute, Maq represents the solute dissolved in the aqueous phase (mol/mL) and Mresin corresponds to the solute adsorbed into the ionic-exchange resin (mol/g). The ratio between the adsorbed solute and the solute that remained in solution is normally described as the distribution coefficient KD (Equation 1).

K D = [M] resin [M] aq Equation 1
In this equation, [M]aq represents the concentration of all dissolved species (mol/mL) and [M]resin represents the concentration of all adsorbed species of M (mol/g). The KD value is very sensitive to the chemical conditions and special care should be taken to the pH of the solution, concentration of solute, the ionic strength of the media, number of exchanging sites and temperature.

To achieve the stability constants, batch experiments at different temperatures using an anionic-exchange resin in the absence or presence of calcium were performed followed by the fitting of a geochemical model to the experimental data. The details are provided below.

M aq ⇌ M resin Reaction 10

In this reaction, M is the solute, Maq represents the solute dissolved in the aqueous phase (mol/mL) and Mresin corresponds to the solute adsorbed into the ionic-exchange resin (mol/g). The ratio between the adsorbed solute and the solute that remained in solution is normally described as the distribution coefficient KD (Equation 1).

K D = [M] resin [M] aq Equation 2
In this equation, [M]aq represents the concentration of all dissolved species (mol/mL) and [M]resin represents the concentration of all adsorbed species of M (mol/g). The KD value is very sensitive to the chemical conditions and special care should be taken to the pH of the solution, concentration of solute, the ionic strength of the media, number of exchanging sites and temperature.

To achieve the stability constants, batch experiments at different temperatures using an anionic-exchange resin in the absence or presence of calcium were performed followed by the fitting of a geochemical model to the experimental data. The details are provided below.

Experimental approach

The distribution of U(VI) between an anionic-exchange resin and the aqueous phase was monitored as a function of different parameters to study the stability of the Ca-U(VI)-CO3 complexes. Further details of the materials and a description of the different methods used in the present study are provided in Appendix A.

Materials and Solutions

The anionic-exchange resin used in the experiments is the gel-type Dowex 1x2 strong base anion exchange resin, stable up to 99ºC. In contrast, the resin used by Dong and Brooks (2006) (Amberlite IRA-410) is not stable at temperature above 40 ºC, and thus, it is not appropriate to study the effect of temperature in the Ca-U(VI)-CO3 system.

The background electrolyte used in the experiments was 0.1 M NaNO3, and initial U(VI) concentrations were in the order of 5•10 -7 M (except when stated).

Batch experiments

Sorption batch experiments were carried out at different temperatures by following, as abovementioned, the experimental procedure proposed by Dong and Brooks (2006).

Before performing these experiments, different types of preliminary tests were performed in order to confirm the correct performance of the experimental set-up. The preliminary tests included:

-To verify if CO2(g) is in equilibrium with the solution at the different temperatures and pressures studied -To test if sampling procedures affect the temperature of the solution in a significant way -To check if there is a significant leaching of organic material from the plastic tubes at the different temperatures studied

After the preliminary experiments were completed, the final experimental set-up used is shown in Figure 6. The experimental set-up consisted in firstly degassing ultrapure water and weight all samples at room conditions and transfer it into the glove-box with controlled CO2(g) (whose concentration depends on the temperature studied). Once inside the glove-box, the stock solutions of calcium and sodium nitrate, sodium carbonate and uranium nitrate solution were prepared and added to each tube. The tubes were sealed with Teflon tape, closed and transferred outside of the glove-box. The samples were then conditioned in vacuum (20 mbar) using an aluminium-foil bag, transferred to an incubator coupled with a reciprocate shaker and left in equilibrium at the temperature of interest.

Once the samples were equilibrated (in this case, after 5 days of contact), the aluminium-foil bag was opened, and the tubes were introduced in an isolated Teflon/aluminium-box, especially prepared to maintain the temperature of the tubes, followed by rapid transfer into the glove-box. In the glovebox, the sampling was done, and the pH was measured at the temperature of interest. The pH meter was previously calibrated with pH 4.0 and pH 7.0 buffer standards at the temperature of interest by using a duo dry-bath system. After cooling the samples at room temperature, the alkalinity of the remaining samples was analysed.

Additional experimental tests were also performed to ensure the viability of the Dowex 1x2 resin for the sorption of U(VI), the initial concentration of U(VI) in solution and the influence of the electrolyte background. These experimental tests were performed in atmospheric conditions at 20ºC, following the procedure proposed by Dong and Brooks (2006).

Two types of batch experiments were performed in controlled atmosphere. These were performed at constant ionic strength, pH, carbonate content and initial concentration of U(VI):

Experiment Type I: Batch experiments at different temperatures in the absence of calcium as a function of the solid-liquid ratio (1-25 g/L). The objective of these experiments is to establish the distribution coefficient of U(VI) in the absence of calcium. Flávia Maia | IMT Atlantique Nantes

Experiment Type II: Batch experiments at different temperatures as a function of the concentration of calcium (0-5 mM) at fixed solid-liquid ratio of 2.5 g/L. The upper limit of calcium concentration was calculated in order to avoid calcite oversaturation in the experiments. The objective of these experiments is to study the stability and the effect of temperature on Ca-U(VI)-CO3 complexes. The different experiments performed are summarized in Table 7. In Callovo-Oxfordian formation conditions (section 1.1.1.1), the increase of temperature is expected to lead to changes in the pH and on the CO2(g) equilibrium of the system (see a description in Tournassat et al., 2008). The theoretical pH and % CO2(g) of each experiment was established taking this effect into account.

For each experimental set, two series of blank samples were used to control the uranyl sorption in the tubes and filters and to control for the possible precipitation of uranium. Further details on these experiments are provided in Appendix A.7.

Conceptual model

The geochemical model approach has been defined to simulate the ionic exchange processes between the resin and the anionic species of U(VI) formed in a solution rich in carbonates, in the absence or presence of calcium, at different temperatures. The calculations aim to determine the stability constants of the aqueous Ca-U(VI)-CO3 species based on the distribution coefficient of U(VI) after equilibration. The solutions were defined according to the equilibrium conditions of batch experiments as a function of the concentration of calcium in solution.

In the conceptual model, each batch is considered as a single system in which each solution is assumed to be in equilibrium with the resin.

In this specific case, the resin is an anionic-exchange resin in the nitrate form (XNO3) and the amount of exchange sites in the resin is calculated based on the density and the exchange capacity of the resin, as described in Appendix B.1.

Different anions present in the solution can exchange with the nitrate in the resin. The overall chemical reactions that may be involved in the exchanger are defined according to the aqueous speciation as follows:

4XNO 3 -+ UO 2 (CO 3 ) 3 4- log 10 K X 1 ⇔ X 4 UO 2 (CO 3 ) 3 + 4NO 3 - Reaction 11 2XNO 3 -+ UO 2 (CO 3 ) 2 2- log 10 K X 2 ⇔ X 2 UO 2 (CO 3 ) 2 + 2NO 3 - Reaction 12 2XNO 3 -+ CaUO 2 (CO 3 ) 3 2- log 10 K X 3 ⇔ X 2 CaUO 2 (CO 3 ) 3 + 2NO 3 - Reaction 13 XNO 3 -+HCO 3 - log 10 K X a ⇔ XHCO 3 + NO 3 - Reaction 14 2XNO 3 -+CO 3 2- log 10 K X b ⇔ X 2 CO 3 + 2NO 3 - Reaction 15
From the reactions above, only Reaction 11 has been considered for the present system. This reaction corresponds to the exchange of UO2(CO3)3 4-, a highly charged species. Thus, its sorption is expected to be more relevant than the sorption of other less charged species, for example UO2(CO3)2 2-, and as such Reaction 12 will not be taken into consideration. According to Dong and Brooks (2006), sorption of CaUO2(CO3)3 2-(Reaction 13) on the resin is expected to be minor. Sorption of Ca2UO2(CO3)3(aq), which is a non-charged species, is expected to be negligible. Finally, the sorption of HCO3 -or CO3 2- (Reaction 14 and Reaction 15, respectively) is considered to play a minor role and will not be taken into account in the calculations.

The concentration of carbonates in the model is controlled by the pCO2 of the system, when no information of its concentration is known. When carbonate concentration is known, the pCO2 of the system will be allowed to be free, where the measured alkalinity and pH will be the main agents that control carbonate concentration.

The conceptual model will be tested by taking into account the experimental data from Dong and Brooks (2006). In this exercise, atmospheric pCO2 will be considered since the concentration of carbonates is known to be in equilibrium with atmospheric pCO2.

Numerical approach

All simulations were conducted by Phreeqc code version 3.1.4 [START_REF] Parkhurst | Description of Input and Examples for PHREEQC Version 3 -A Computer Program for Speciation, Batch-Reaction, One-Dimensional Transport, and Inverse Geochemical Calculations[END_REF] which includes chemical equilibrium reactions and gas-aqueous phase equilibria.

The determination of the stability constants of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) (log10β113 and log10β213, respectively) considering the experimental distribution coefficient of U(VI) at different temperatures was done using Parameter ESTimation code PEST [START_REF] Doherty | PEST Model-Independent Parameter Estimation User Manual[END_REF].

PEST is an opensource, public-domain software suite that allows model-independent parameter estimation and parameter/predictive-uncertainty analysis. Subsequently, PEST was coupled to the batch version of Phreeqc.

A benchmark was also performed by using a complementary approach with the aim of fitting the model to the experimental data. This was achieved by designing a matrix with Microsoft Excel © where two variables (log10β113 and log10β213) are defined simultaneously. The PhreeqC input file is defined to run all combination possibilities within a reasonable and equally spaced values range. The sum of squared differences (SSD) and then calculated in Microsoft Excel © between the model and the experimental data. The OriginLab© software was used to represent graphically the variations of the SSD values as a function of both variables. The lower SSD value is taken as the best approach and it is within the darkest zone of the 2D mapping. To reduce the uncertainties of the variables, the previous routine can be applied to a more limited range of values.

All calculations were performed using ThermoChimie Database v9b with a Davies approach for ionic strength correction, unless otherwise mentioned. The database already contains the stability constants for the CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) complexes, based on the work by Dong and Brooks (2006). These data have not been used in the calculations as one of the objectives of the present work is to confirm those values. On the other hand, the thermodynamic data for UO2(CO3)2 2-and UO2(CO3)3 4-is also included in this database, where the stability constants and enthalpy values originate from the work of Guillaumont et al. (2003). These U(VI)-CO3 complexes are the most relevant U(VI) complexes expected to be formed in the absence of calcium under the conditions of the study, and that data has been included in the calculations.

An extensive listing of aqueous and solid species included in the database relevant for the study is provided in Appendix E.

Application of the conceptual model on Dong and Brooks (2006) experimental data

According to the conceptual model proposed in the present study, the distribution coefficient of U(VI) in the absence or presence of calcium is calculated based on the sum of its speciation distribution between the aqueous and resin phase. The model approach that we applied takes into consideration the isotherm of U(VI) onto the anionic exchange resin in the absence of calcium in solution to determine the stability constant of the anionic-exchange reaction (Reaction 11). Furthermore, a model approach using the stability constant determined previously and the experimental KD of U(VI) as a function of calcium in solution is applied to determine the stability constants of the Ca-U(VI)-CO3 ternary complexes. The results of both model approaches are shown in Figure 7. All constants were fitted by PEST.

The model approach in both cases reproduces well the experimental data, both the U(VI) sorbed onto the resin (Figure 7a) and the KD of U(VI) as a function of calcium (Figure 7b). The results of the fitting approach indicate that the log10 βº113 and log10 βº213 for the Reaction 1 and Reaction 2 are respectively 27.29 (± 0.07) and 30.02 (± 0.63). These results differ slightly from those proposed by Dong and Brooks (2006) in which log10 βº113 = 27.18 (± 0.07) and log10 βº213 = 30.7 (± 0.06), although they are in the same range. The most important issue that should be pointed out is the uncertainty of the resultant stability constant. Special concern should be taken on the calculation of the uncertainty of log10β213, which seems to be very small in the case of the work by Dong and Brooks (2006). In conclusion, the conceptual model here presented can be validated based on the application of the model approach on the data of Dong and Brooks (2006). However, with an increase in the temperature of the system some considerations must be taken:

Not all species or solids have enthalpy of reaction values, which will be considered zero for the modelling calculations. This assumption must be taken carefully, since the behaviour of some species might change with an increase in temperature.

Results

Experimental tests

Extensive details of the experimental test results are available in Appendix C.1.

Controlling test on temperature and pCO2

Different preliminary tests were performed to guarantee the correct performance of the experimental set-up (Figure 6). A summary of the experimental tests and results obtained are described in Table 8.

Challenge Task Actions

Equilibrium with CO2(g) is difficult to reach in solutions at T > 25ºC

Study carbonates concentration at the different temperatures and pressures studied.

pH will be measured inside the glovebox at the temperature of the experiments. Alkalinity will be measured at room temperature inside the glovebox. Temperature may change during sampling procedures Verify if the temperature is maintained when sampling An isolation box will be used to control the temperature Degradation of PPCO tubes at high temperatures Verify by TOC analysis the concentration of organic carbon released in time at different temperatures Washing 3 times the tubes with ultra-pure water at respective temperature before each batch experiment. Some insights from the results of the CO2(g) equilibrium experiment (first test) will be provided here, since they are essential for the interpretation and modelling of the experimental data.

The equilibrium of CO2(g) with the solution may be difficult to achieve, and can be a challenge at temperatures higher than 25ºC [START_REF] Giffaut | Influence Des Ions Chlorure Sur La Chimie Des Actinides Effects de La Raiolyse et de La Température[END_REF]. In order to evaluate this, alkalinity measurements were done following the procedure described in Appendix A.3. The results are shown in Table 9. The results show that the equilibrium between the chemical system and the pCO2 is not reached during the time of the experiment. The major differences are observed at higher temperatures (60 and 80ºC). To minimise the uncertainty of the calculated pCO2 values, it is crucial to measure the pH and the alkalinity for all experimental cases. These parameters will be further used in the simulation/fitting to obtain correct interpretation and modelling data. 10 -2.1 10 -2.2 10 -2.0 10 -2.2 10 -1.5 10 -1.8 10 -1.3 10 -1.7

Controlling test on the initial concentration of U(VI) and electrolyte background

Besides the lack of information on the control of carbonates in the system, another issue that is almost as disregarded in all published studies is the lack of evaluation of the possible precipitation of U(VI)bearing minerals, such as clarkeite and calcite. Bernhard and co-workers (2001) and Dong and Brooks (2006) were the only ones monitoring the concentration of uranium and calcium to check the possible precipitation of mineral phases. In both works, it was reported that no solids were observed. In all experimental data published (see details in Appendix D. 1), the initial concentration of U(VI) in solution (2•10 -5 and 4.7•10 -4 M) is high enough to precipitate some U-mineral phases. Not only uranium plays an important role on the solubility of U-bearing minerals but also the concentration of sodium and calcium are relevant for the solubility of such mineral phases. For example, clarkeite (Na-Uranate) easily precipitates even at low ionic strength solutions (0.01 M) [START_REF] Giammar | Influence of Dissolved Sodium and Cesium on Uranyl Oxide Hydrate Solubility[END_REF]).

Another example is the possible precipitation of becquerelite (Ca-U(VI) hydroxide mineral), which solubility is below 1.5•10 -8 M of U(VI) [START_REF] Casas | Characterization and Dissolution Behavior of a Becquerelite from Shinkolobwe, Zaire[END_REF][START_REF] Rai | Thermodynamics of the UVI)-Ca 2+ -Cl --OH-H2O System: Solubility Product of Becquerelite[END_REF][START_REF] Sandino | Solubility Equilibria in the U(VI)-Ca-K-Cl-H2O System: Transformation Od Schoepite into Becquerelite and Compreignacite[END_REF]. Consequently, using concentration of U(VI) above this value can lead to an oversaturation of becquerelite. However, the possible precipitation of mineral phases should be merely speculative as the formation of these type of phases is not controlled by concentration of U(VI) in solution but also by its kinetics. For this reason, assuming the precipitation of such mineral phases in the geochemical calculations may lead to misinterpretation of the experimental data because the equilibrium is always assumed.

Predictive calculations using Thermochimie Database v9b (Giffaut et al., 2014) performed taking the experimental data reported by Dong and Brooks (2006) and Bernhard et al. (2001) show that calcite, clarkeite and becquerilite seem to be oversaturated despite the observations reported by the authors. Dong and Brooks (2006) indicated that calcite may precipitate at concentrations of calcium above 3 mM (as also observed in these predictive calculations); although in case of the U(VI)-bearing minerals would be undersaturated under their experimental conditions (Figure 8a). On the other hand, Bernhard and co-workers (2001) reported that even after one year no precipitation or other chemical changes were detected. The same results were observed using the experimental data reported by Endrizzi and Rao (2014) (Figure 8b). Details of all modelling results of the experimental conditions for all works described in Table 4 are provided in Appendix D.1.

Except for the works of Endrizzi and Rao (2014) and Lee and Yun (2013), most published work mentions that no mineral precipitation was observed, and a careful filtration of the samples was taken. Even though the authors did not mention the possibility of precipitation it does not mean that they did not considered this issue. Most probably they did not observe any perturbation of the system and realised that it was not worth mentioning it.

It should also be noted that the solubility of the U(VI) solids is not only dependent on the initial concentration of U(VI), but also on other chemical parameters such as pH, temperature and concentration of carbonates. Following the discussion above, the batch experiments carried out in the present study were performed using 5•10 -7 M of initial concentration of U(VI) (this means two orders of magnitude lower than reported by Dong and Brooks work (2006)). According to geochemical calculations, no U(VI) phases are expected to be oversaturated, at least at temperatures below 25ºC, under the experimental conditions envisaged for the present work. If they are, this will be taken into account on the calculations of the stability constants.

The composition of the electrolyte background is a relevant parameter to consider in this kind of studies. It is well known, that the composition of electrolyte media is important due to its interaction with the uranium in solution. For this reason, a less complexing electrolyte media should be used. Except for Dong and Brooks (2006) and Endrizzi and Rao (2014), that used NaNO3 and NaCl (respectively) as electrolyte background, other experimental works used sodium perchlorate, which is a less complexing agent in its relation to uranium.

For this reason, to test the impact of taking sodium perchlorate as electrolyte background instead of sodium nitrate as proposed by Dong and Brooks (2006), a batch experimental test was performed with 0.1 M NaClO4.

Special care must be taken upon using NaClO4 as electrolyte background in experiments, as it is well known that pH measurements under perchlorate conditions can affect the pH electrode due to precipitation of KClO4 at the junction of the electrode. Moreover, the pH electrode used in this experimental test is very sensitive to reference electrode solution changes, thus the reference solution was not changed. The pH measurements during this test were performed as quickly as possible to avoid the precipitation of KClO4. Another issue that should be considered, besides the impact of the perchlorate solution in pH measurements at room temperature, concerns the increase of the temperature in the system. This increase may lead to an increase in the solubility of the salt solution, which will also interfere with pH measurements.

Nonetheless, the ionic exchange batch experiments using sodium perchlorate were performed at room temperature and atmospheric conditions for practical reasons. The resin was previously changed for perchlorate form (as for the nitrate form, see Appendix A.6). The pH of the equilibrium was measured at the end of the experiment and found to be an average of 7.8 (± 0.1). The details of the experimental results are provided in Appendix C.3. The results of the concentration of U(VI) in solution are shown in Figure 9. The sorption of U(VI) in the anionic exchange resin is almost negligible, showing a strong selectivity coefficient of the perchlorate with respect to the anionic species of U(VI) present in solution. The same effect was observed in [START_REF] Gu | Sorption and Desorption of Perchlorate and U(VI) by Strong-Base Anion-Exchange Resins[END_REF] when different types of anion-exchange resins were tested with both U(VI) and perchlorate. The authors show that the resin's functional group plays an important role on the selectivity of the anionic species, and that approximately 34 % of perchlorate can be desorbed from gel-type resins (such as Dowex 1x2) after rinsed with HCl diluted acid.

In the present study, the perchlorate is hardly desorbed from the resin. Consequently, the distribution coefficient of U(VI) is constant and the formation of Ca-U(VI)-CO3 ternary complexes cannot be studied. As a result, the use of perchlorate as counter-ion on the anionic-exchange resin and as electrolyte media was dismissed and all ionic exchange batch experiments were performed using sodium nitrate as performed in Dong and Brooks (2006).

Determination of the stability constants in Ca-U(VI)-CO 3 -H 2 O systems at different temperatures

After validation of the proposed experimental procedure, we studied the aqueous behaviour of U(VI) under a controlled atmosphere at different temperatures. The study consists in changing the pCO2 of the system according to the temperature of interest (20, 40, 60 or 80 °C). The experimental conditions try to reproduce as closely as possible the chemical system of COx formation (section 1.1.1). The resultant KD further allows the determination of thermodynamic properties of the Ca-U(VI)-CO3 aqueous species. The details of the experimental results are provided in Appendix C.3, and the modelling results and the input files can be found in Appendix D. 2 and F, respectively.

In the absence of calcium

Batch experiments in the absence of calcium (Type I, Table 7) were carried out to determine the distribution coefficient of U(VI) in the absence of Ca-U(VI)-CO3 complexes.

U(VI) initial concentrations in the experiments are ca. 5•10 -7 M. No precipitation was visible in the tubes, neither before nor after filtration and this was confirmed by measure the blanks and afterwards by Phreeqc calculations. The alkalinity measured in the blank samples led to a total concentration of dissolved inorganic carbon in solution of around 2 meq/L.

The experiments were performed as a function of the solid-liquid ratio (1-25 g/L). An increase of the solid-liquid ratio promotes a pH decrease (Figure 10) probably related to changes on the gas/solution volume. This has been taken into account in the modelling calculations. As explained previously (section 2.2.1.1), at temperatures higher than 25 ºC it is difficult to reach an equilibrium between the solution and the CO2(g) atmosphere. Thus, the measured alkalinity values shown in Table 9 have been used for the interpretation of the results obtained (Figure 11). The results obtained in Type I experiments show that the concentration of U(VI) in solution (symbols in Figure 11) decreases as the amount of resin increases. The effect of temperature on the concentration of equilibrium of U(VI) is directly controlled by the thermodynamic properties of water and simultaneously by the thermodynamic properties of the speciation of U(VI) (Appendix E), assuming that there are no changes on resin properties. The speciation of U(VI) in the presence of carbonates in temperatures between 20 and 80 ºC is controlled mainly by UO2(CO3)3 4-and UO2(CO3)3 2-(Figure 12). At 80 ºC, the hydrolysis of the UO2(OH)3 -and UO2(OH)2(aq) are expected but can be considered negligible. These changes of the speciation of U(VI) in solution, due to the changes in temperature, may lead to changes of the Log10KX. At 20 ºC, the system can be simple, since there is only one predominant species that may compete in adsorption to the resin. However, with increasing temperature the complexity of the chemical system increases, and two aqueous species may compete for the adsorption of U(VI). The experiment proposed here cannot distinguish which aqueous species will control the sorption of U(VI) onto the resin, therefore it is not possible to determine two different Log10KX at the same time (Reaction 11 and Reaction 12). Thus, the experimental data have been modelled using PEST in order to determine a constant for UO2(CO3)3 4-sorption (Reaction 11) at each temperature studied. It should be pointed out that these experiments have been performed to quantify the U(VI) adsorbed to the resin in the absence of calcium, which will be used as a reference state when calcium is present and the formation of Ca-U(VI)-CO3 ternary complexes are favoured. In fact, description of the real mechanism of the anionic-exchange processes is not of concern to this work. Thus, as a simplification of the geochemical model, it will be assumed that the UO2(CO3)3 4-will be the only species adsorbed to the anionic-exchange resin. This is assumed as these species are present in all cases and also have high anionic charge, more likely to be adsorbed to the resin, in contrast to the lower anionic charge species.

The agreement between the experimental data and the calculated values at 20, 40 and 60 ºC is excellent (solid-lines in Figure 11). In contrast, the fit between the estimated and experimental observations at 80 ºC and 5 % of CO2(g) is less satisfactory. Although some deviation is noticed in the experiments at 80 ºC, the agreement between the model and the experiments at S/L = 2.5 g/L (which is the ratio that will be used afterwards in the experiments in the presence of calcium, see further Section 2.3.1.2.2) is excellent.

The results show that there is a clear effect of temperature on the anionic-exchange resin, following a linear trend, in which the complexation of UO2(CO3)3 4-onto the anionic-exchange resin is favoured with a temperature increase (Figure 13). The log10 KX1 values for Reaction 11 obtained from the fitting procedure at each temperature are provided in Table 10. The dependency of the log10 KX1 values with temperature is shown in Figure 13. The results indicate that the dependence of the stability constants with

1 T 0 - 1 T is a straight line.
The slope of this line can be correlated with an "apparent" enthalpy of reaction for the process represented in Reaction 11 using the Van't Hoff equation (Equation 3), assuming that the standard heat capacity is constant.

log 10 K °(T) = log 10 K °(T 0 ) + ∆H r °(T 0 ) Rln(10) ( 1 T 0 - 1 T ) Equation 3
In Equation 3, R is the gas constant (8.31451 JK -1 •mol -1 ).

Thus, using the fitted data for the log10 KºX1 (25 ºC) values at different temperatures and applying the Equation 3, it is possible to calculate the standard stability constant and an "apparent enthalpy" (ΔrH app ) value for Reaction 11. The resultant log10 KºX1 is 1.1 (± 0.06) and the ΔrH app is 21.2 (± 3.3) kJ/mol.

From the results of the batch experiments of anionic exchange of U(VI) in the absence of calcium it is possible to: quantify the U(VI) sorption to the resin; verify that the simplified conceptual model proposed in the present work is able to explain the experimental results obtained and; determine the standard stability constant and the "apparent" enthalpy values for the Reaction 11 to be used in modelling calculations in the presence of calcium.

In the presence of calcium

We performed batch experiments of sorption of U(VI) on resin as a function of calcium concentration (Type II, see Table 7) at different temperatures in order to study the formation of Ca-U(VI)-CO3 complexes.

The initial concentration of U(VI) was set as ca. 5•10 -7 M. As before, no precipitation was visible in the tubes neither before nor after filtration. The average pH and alkalinity measured in the blank samples are shown in Table 11. There is good agreement between the pH calculated and the one measured (compare values described in Table 11). This means that the equilibrium of pCO2/pH/temperature of the present system is relatively well controlled. The results obtained in Type II experiments show that the concentration of U(VI) in solution (Figure 14) increases as the calcium concentration in the system increases. This behaviour is related to the formation of Ca-U(VI)-CO3 complexes, according to the conceptual model described in Section 2.2.1.2. Contrary to the observed behaviour of the U(VI) in the absence of calcium, the retention of U(VI) in the resin decreases when the temperature increases. In Figure 14, two sets of data can be clearly distinguished: a) data at 20 and 40 ºC and b) data at 60 and 80 ºC. The concentration of uranium in equilibrium is higher at 20 and 40 ºC than at 60 or 80 ºC. The change of concentration can be related to the changes of the speciation of uranyl-carbonate complexes as observed in Figure 12. This means that the formation of UO2(CO3)3 4-and CaUO2(CO3)3 2-is strongly affected by the temperature of the system. The respective stability constants for the Reaction 1 and Reaction 2 are calculated as shown in the following Equation 4.

β x13 = [Ca x UO 2 (CO 3 ) 3 2x-4 ] [Ca 2+ ] x [UO 2 2+ ][CO 3 2-] 3 Equation 4
In Equation 4, x = 1 or 2. The log10 KºX1 and the "apparent" enthalpy determined previously were used to fit the experimental data of concentration of U(VI) as a function of calcium at the different temperatures. The best fit line followed a non-linear model since two species CaUO2(CO3)3 2-and Ca2UO2(CO3)3 are formed. The results obtained in the fitting of the experimental data are provided in the following sections.

Fitting with PEST

In a first approach, the fitting of the stability constants has been performed with PEST coupled to Phreeqc. The results of the model approach for each experimental data are shown in Figure 15 and Figure 16. The agreement between the model approach and the experimental data is excellent for each experimental set. The log10 β113 and log10 β213 for each temperature obtained from the fitting are provided in Table 12. The details of the fitting are provided in Appendix D.2 and F.

The fit of the model approach with increasing temperature and respective pCO2 of the chemical system, with the experimental data is good (solid-lines in Figure 15). The model approach also reproduces well the KD of U(VI) for all experimental cases (solid-lines in Figure 16) except at 40 ºC, since the calculated values are slightly overestimated at low concentrations of calcium with respect to the experimental data (Figure 15b and Figure 16b). The dependency of temperature on the stability constants of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) is shown in Figure 17. From the model results, it is possible to observe that both constants do not exhibit the same behaviour, as log10 β113 clearly decreases with temperature and log10 β213 is not affected by temperature changes.

Benchmark of fitting using 2D mapping

Even though the use of PEST was successful for the parameterisation of the stability constants, this software can lead to high uncertainty, as the definition of the global minimum in non-linear systems is difficult to achieve. For this reason, a complementary numerical approach was performed by using Phreeqc as a geochemical model software, Excel for the sample standard deviation (SSD) and OriginLab© to evaluate the range of stability constants by 2D mapping.

With this method, the best fit for the stability constants is defined by the difference of SSD by means of colour grades. The lightest colour corresponds to the lower probability of the stability constants values (higher SSD) and the darkest correspond to the higher probability that can occur in a pair of stability constants (lower SSD). The best pair of stability constants lies within the darkest zone of the 2D mapping. The results of 2D mapping for each Type II experiment are shown in Figure 18. The circle-and square-sign represent the optimised value of the stability constants for each fitting method; i.e. the circle-sign corresponds to the best fit determined from matrix defined by Excel coupled with Phreeqc, and the square-sign corresponds to the best fit provided by PEST coupled with Phreeqc. The results shown in Figure 18 indicate that the constants for Ca-U(VI)-CO3 complexes seem to be inversely correlated; for higher values of log10β213, lower values of log10 β113 are calculated. For the experiments at 20 ºC and 40 ºC, 2D mapping indicates that the best fit is strongly dependent on the stability constant of Ca2UO2(CO3)3(aq). This is probably related to the fact that this complex may play a minor role in the experiments at those temperatures (see further Figure 23).

The optimised values of the stability constants for the Ca-U(VI)-CO3 ternary complexes at different temperatures and pCO2 are described in Table 13. A further comparison of the calculated stability constants as a function of temperature using PEST or using 2D mapping are provided in Figure 19. The results obtained with both approaches are in good agreement and, in most cases, they are within the uncertainty range calculated with PEST. Highest deviations are observed in the case of data obtained at 20 and 40 ºC. From the experiments undertaken at different temperatures, we determined the thermodynamic data needed to understand and model the U(VI) behaviour under temperature elevation. One can write the temperature dependence of the equilibrium constants as a function of the enthalpy, since the range of temperatures is small enough that the enthalpy can be assumed constant and assuming that the heat capacity change of reaction is zero at all temperatures (Equation 5). The enthalpy of reaction and the standard stability constants (at 298.15 K) were calculated following Equation 3.

Evaluation of stability constants and enthalpy of reaction

Knowing the standard stability constant and the enthalpy of reaction, we further can determine the Gibbs free energy by its isotherm equation (Equation 5) and the entropy (Equation 6).

𝛥 r G°= -RTlnK°(T 0 ) Equation 5

Δ r G°=Δ r H°-TΔ r S° Equation 6

In Equation 5and Equation 6, R is the gas constant (8.31451 JK -1 •mol -1 ). We obtained the thermodynamic data for both uranyl-calcium-carbonate complexes, based on these calculations. The results summarized in Table 14 are the average values from both the PEST and 2D mapping approaches. As observed in the literature, the stepwise constants determined in the present work agree with the expected trend: the binding constant of Ca 2+ to UO2(CO3)3 4-(Reaction 16) is larger than the Ca 2+ binding constant to CaUO2(CO3)3 2-(Reaction 17).

Ca 2+ + UO 2 (CO 3 ) 3 4-i ⇌ CaUO 2 (CO 3 ) 3 2- log10 Kº = 5.46 Reaction 16 Ca 2+ + CaUO 2 (CO 3 ) 3 2-⇌ Ca 2 UO 2 (CO 3 ) 3 (aq) log10 Kº = 2.39 Reaction 17
The slope of the linear correlation between the stability constants and the temperature indicate that the enthalpy of reaction is different for each uranyl-calcium-carbonate ternary complex. The formation of the anionic species is not favoured with the increase of temperature controlled by an exothermic reaction; whereas the formation of the neutral species is not significantly affected by the changes of temperature. The enthalpy of reaction of the neutral species is slightly positive, although i Note: 𝑈𝑂 2 2+ + 3𝐶𝑂 3 2-⇌ 𝑈𝑂 2 (𝐶𝑂 3 ) 3 4-, LogK 0 = 21.84 according to data selection in Guillaumont, 2003. its uncertainty is very high, so it can be assumed that there is no impact of the temperature on the formation of the Ca2UO2(CO3)3(aq). We obtained the calculations shown in Figure 21 and Figure 22 using the stability constants and enthalpy of reaction provided in Table 14 and observed a good agreement between experimental and calculated data. The uranium aqueous speciation values calculated for the experiments at the different temperatures are shown further in Figure 23. The aqueous speciation results indicate that CaUO2(CO3)3 2-is the main aqueous U(VI) species at all temperatures for Ca concentrations above 1mM. The fraction of Ca2UO2(CO3)3(aq) increases with Ca concentrations and in temperature, but this species is not the main one in any case. This is linked to the higher uncertainty found in the determination of its stability constant.

Comparison with previous literature data

The stability constants for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) obtained as an average of PEST and 2D mapping fitting are log10 βº113 = 27.3 (± 0.2) and log10 βº213 = 29.7 (± 0.3). These values obtained from ionic exchange methods at different temperatures are consistent with the values previously found (Figure 24). The stability constant log10 βº113 agrees with the recent works of Dong and Brooks (2006), Endrizzi and Rao (2014) and Lee and Yun (2013). The values for log10 βº213 are in agreement with the published literature values within the range of uncertainty.

Although the stability constants obtained here for both complexes are in agreement with the published literature values within the range of uncertainty, taking as example the Callovo-Oxfordian geochemical conditions (pH = 7.2, pCO2 = 10 -2.1 atm and [U(VI)] = 10 -8 M) we can evaluate the impact of the stability constants of those complexes. Comparing the stability constants proposed in the present work with those reported by Dong and Brooks (2006), we can observe significant changes on the speciation of U(VI) as a function of calcium in solution (Figure 25). The predictive modelling results at 25 ºC indicate that CaUO2(CO3)3 2-plays an important role on the U(VI) speciation with the increase of the concentration of calcium in solution at the expense of Ca2UO2(CO3)3(aq). Even if the change in calculated speciation is significant, the impact on uranium solubility is more limited, as shown below. The solid exerting a solubility control of uranium in clay porewater conditions (Callovo-Oxfordian porewater, Table 3, see Section1.1.1.1) at 25 ºC is expected to be UO2•2H2O(am). The calculated aqueous uranium speciation under Callovo-Oxfordian conditions in equilibrium with UO2•2H2O(am) is shown in Table 15.

Two different calculations, using two different sets for the stability constants of Ca-U(VI)-CO3 complexes, have been performed: using the stability constants form Dong and Brooks (2006); and using the stability constants determined in the present work. As seen in Table 15, even if the calculated U(VI) speciation is very different, the calculated uranium solubility is similar in both cases.

The enthalpy of reaction for CaUO2(CO3)3 2-determined in the present work is almost two times smaller than the one reported by Endrizzi and Rao (2014). The most important difference is observed in the case of Ca2UO2(CO3)3(aq) thermodynamic data, although it must be kept in mind that the uncertainty associated with the formation of this complex is also high (see the discussion above). One reason behind those differences may be the possible uranium precipitation in the experiments by Endrizzi and Rao (2014), due to the high uranium initial concentration needed in the microcalorimetry experiments. The overall thermodynamic data of the Ca2UO2(CO3)3(aq) and CaUO2(CO3)3 2-from the present work and the data of Endrizzi and Rao (2014) can be compared in the following Table 16. Finally, we tested if the previous literature data could be used to explain the experimental measurements obtained in the present work. In this context, the results obtained in the ionic exchange experiments in the presence of calcium have been modelled using the thermodynamic data for Ca-U(VI)-CO3 complexes, available in Endrizzi and Rao (2014), since both stability constants and enthalpy data are reported.

The comparison of modelling with the thermodynamic data of the present work and with data from Endrizzi and Rao ( 2014) is shown in Figure 26. The modelling results indicate that the thermodynamic data for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) using the data from Endrizzi and Rao (2014) provide reasonable results at low temperatures, but a higher deviation is observed when the temperature increases.

Summary

In conclusion, the experimental method proposed here, in which we determined the stability constants of the Ca-U(VI)-CO3 ternary complexes at different temperatures, and consequently their thermodynamic parameters, was successful.

The anionic-exchange experimental method performed in the present study was adapted from the one proposed by Dong and Brooks (2006), where the main modifications were attributed to the changes of anionic-exchange resin more stable in temperature, and the initial concentration of U(VI). Furthermore, preliminary experiments were performed to guarantee the correct performance of the experimental set-up and to evaluate the equilibrium between CO2(g) and the solutions.

The main issue found for this experimental method was the difficulty in controlling the changes of pH/carbonates of the system in temperature during the batch experiments. To achieve the control of all parameters in this type of systems was challenging. Initially, the experimental tests show that the pH and alkalinity (control of the carbonate concentration) were slightly higher than expected. Besides these less accurate results during the anionic-exchange batch experiments, the measured pH and alkalinity of the solution were in good agreement with those from the theoretical data. Flávia Maia | IMT Atlantique Nantes

The batch experiments in the absence of calcium aiming to quantify the sorption of uranyl-carbonate complexes on the resin at different temperatures (20, 40, 60 and 80 ºC) and controlled atmospheres were successfully accomplished. Moreover, we also performed batch experiments in the presence of calcium adapted from Dong and Brooks (2006) to study the formation of Ca-U(VI)-CO3 complexes at different temperatures (20, 40, 60 and 80 ºC) and controlled atmospheres.

We determined the stability constants of Ca-U(VI)-CO3 complexes by fitting a geochemical model, previously evaluated and validated using the Dong and Brooks (2006) experimental data, to the experimental data obtained from the anionic batch experiments performed at different temperatures.

The fitting was performed by two different parameterisation/optimisation software (PEST and Ex-cel©). The results show that both aqueous complexes studied here show different behaviours in temperature. The formation of the anionic complex (CaUO2(CO3)3 2-) is not favoured with the increase of temperature, though the formation of the neutral complex (Ca2UO2(CO3)3(aq)) is independent of the changes in temperature (Figure 20).

From these data, the standard stability constants for CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) were obtained and are consistent with those previously published, which provides confidence on the methodology used and the conceptual model proposed. Furthermore, we also determined the enthalpy data for the formation of Ca-U(VI)-CO3 complexes (Table 14). The enthalpy of reaction of the formation of CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq) are -27.4 ± 8.2 and 0 ± 1.8 kJ/mol, respectively. The ΔrHº of the anionic complex is negative as proposed by Endrizzi and Rao (2014), although it is two times lower. However, for the neutral complex, the ΔrHº contradicts the value proposed by Endrizzi and Rao (2014) (-47 kJ/mol). Besides the high uncertainty of ΔrHº of the Ca2UO2(CO3)(aq) reaction, there is no impact of the temperature on its formation, in which we can assume that the ΔH is zero.

The other thermodynamic parameters of CaUO2(CO3)3 2-and Ca2UO2(CO3)(aq) formation (entropy and Gibbs free energy) are in good agreement with those published by Endrizzi and Rao (2014) (Table 16).

Chapter 3 Behaviour of Uranium in Callovo-

Oxfordian Formation at 20 ºC

Introduction

The development of reliable predictive sorption models is essential to accurately evaluate the longterm fate of radionuclides in real geochemical conditions. The aim of this chapter is to describe the retention mechanisms of U(VI) in Callovo-Oxfordian formation at 20 ºC. The model here proposed for the retention mechanisms of U(VI) will be further used as a basis for studying the temperature effect on the sorption of U(VI) on Callovo-Oxfordian systems at different temperatures (see further Chapter 4).

One of the difficulties of the present work lies on the complexity of the Callovo-Oxfordian claystone (carbonated clayed facies) mineral assemblage (see Table 1, [START_REF] Gaucher | ANDRA Underground Research Laboratory: Interpretation of the Mineralogical and Geochemical Data Acquired in the Callovian-Oxfordian Formation by Investigative Drilling[END_REF]. In general, the clay fraction of COx clayctone corresponds to 20-50 wt% of the total mass of the formation, where illite and interstratified illite/smectite (disordered I/SR0 and ordered I/SR1) are the main components. This argillaceous rock is well known for its high content of carbonates, such as calcite and dolomite that can correspond to 20 to 40 wt%, responsible for buffering the system from neutral to weakly alkaline conditions. The minor content of pyrite present in the COx formation (< 1 wt%) allows the maintenance of the redox potential as reductive [START_REF] Vinsot | In Situ Characterization of the Callovo-Oxfordian Pore Water Composition[END_REF][START_REF] Vinsot | O2 Consumption and CO2 Production at Callovian-Oxfordian Rock Surfaces[END_REF]. The presence of clay minerals, in addition to a neutral pH environment and reducing aqueous conditions, are relevant parameters for the safety of nuclear waste, since all favour a high sorption of radionuclides and a low solubility of most radionuclides of interest, such as uranium, by limiting radionuclide migration through the geosphere [START_REF] Bruno | The Applicability and Limitations of Thermodynamic Geochemical Models to Simulate Trace Element Behaviour in Natural Waters. Lessons Learned from Natural Analogue Studies[END_REF] The present study aims to answer the following questions: what mineral phases contribute for the retention of metal/radionuclides ions and which retention mechanisms should be considered in the model to describe the retention of U(VI) under Callovo-Oxfordian conditions?

We propose a three-step methodology:

1) A description of the main phases that could be involved in the retention processes;

2) A literature review to select and define a retention model and to identify possible missing parameters with respect to Callovo-Oxfordian formation conditions; 3) The definition of the experimental parameters to be measured during the laboratory experiments will be discussed in light of the literature review.

Background

In the following section, we provide a brief description of the main solid phases that contribute for the retention of radionuclides in Callovo-Oxfordian claystone systems as well as the mechanisms of retention involved at the water-minerals interface and we will discuss how these can be interpolated for natural systems.

What are clay minerals?

The clay minerals are hydrated phyllosilicates composed by tetrahedral and octahedral sheets separated or not by interlayer cation and/or water molecules. They are classified according to: (1) the thickness between the tetrahedral and octahedral sheets;

(2) the anisotropy of the layers or particles;

(3) the existence of several types of surfaces (planar or basal and edges); (4) the ability to which the external (and internal) surface can be modified by e.g. adsorption and ionic exchange; (5) plasticity; and (6) hardening by heat [START_REF] Brigatti | Chapter 2 Structures and Mineralogy of Clay Minerals[END_REF].

The clay minerals contain continuous tetrahedral (T) and octahedral (O) sheets. Each tetrahedron consists of a cation (Si 4+ , Al 3+ and Fe 3+ ) coordinated to four oxygen atoms (Figure 27b) and linked to adjacent tetrahedral through shared basal oxygen atoms. The octahedron consists of a cation (Al 3+ , Fe 3+ , Mg 2+ and Fe 2+ ) coordinated to six oxygen atoms (Figure 27a,c) and linked to the next octahedral sheet by sharing edges. The combination of several tetrahedra across basal oxygen atoms leads to the formation of a basal surface, while for the octahedral sheets the individual octahedra are connected by sharing edges. When Al 3+ is the central cation, only two of every three octahedra are occupied to maintain electrical neutrality resulting in a so-called dioctahedral structure (Figure 27d). On the other hand, when Mg 2+ is the cation, all three octahedra are filled forming a so-called trioctahedral structure (Figure 27e).

The combination of one tetrahedral sheet and one octahedral sheet results in a 1:1 (TO) layer structure (Figure 28a). The edges of the tetrahedral sheets project into a hydroxyl plane of the octahedral sheet, replacing 2/3 of the hydroxyl ions and forming one basal plane of oxygen atoms from the tetrahedral sheet while the other plane is made of hydroxyls from the octahedral sheet [START_REF] Bergaya | Uranyl (VI) Carbonate Complex Formation: Validation of the Ca2UO2(CO3)2(aq.) Species[END_REF]. The combination of one octahedral sheet sandwiched between two tetrahedral sheet results in a 2:1 (TOT) layer structure (Figure 28b). Here both basal planes of individual layers are composed by oxygen atoms (belonging to the tetrahedral sheets). Isomorphous substitutions may occur in these layer structures. The isomorphous substitutions are characterised by replacing the Si 4+ in tetrahedral positions by cations of a similar size and coordination but of lower valency, or/and through substitutions of Al 3+ in octahedral positions. As a result, the layer structure acquires permanent negative charge. The isomorphous substitutions are typical of 2:1 layer structures in which the net layer charge is balanced by hydrated or non-hydrated-cations occupying interlayer positions. When the interlayer space is filled with a hydroxide sheet a 2:1:1 layer structure is formed. Based on these parameters, we can distinguish different groups of clay minerals.

In So far, we described the clay mineral as a single unit cell. When considering all unit cells, the planar phyllosilicates are formed by a finite number of either 1:1 or 2:1 type layer structures that are continuous along the layer planes and stacked perpendicularly. This kind of arrangement leads to the possibility of layer displacements by rotation and/or translation, which provides several polytypes or structures with different layer stacking sequences. The packing of the different types of layers can be randomly or regularly interstratified in a so-called mixed-layer structural arrangement. The interstratified clay minerals can be classified based on their stacking as ordered, random or partially ordered.

In Figure 29, we show a schematic example of the stacking of interstratified clay minerals.

What are the retention processes?

The retention processes can be classified as reversible or irreversible. In the reversible retention processes, there are sorption processes such for example: ion exchange and surface complexation. On the other hand, irreversible retention is associated generally with immobilisation processes (in which in some cases can be reversible in very slow equilibrium processes) such as surface precipitation, coprecipitation, surface substitution or diffusion on the minerals (NEA, 2002),.

In real geochemical conditions, all retention processes should be included in prediction models. However, due to the complexity of most natural systems, simplifications must be made by neglecting, in some cases, some of these processes. For instance, in predictive models most irreversible processes are not considered due to a lack of relevant data, and the presence of uncertainties associated to the slowness and complexity of such processes [START_REF] Nea | Radionuclide Retention in Geologic Media[END_REF]. The understanding of irreversible processes has been improved due to the recent development of high-quality laboratory surface-specific analyses such as Surface Nuclear Magnetic Resonance (NMR) or Extended X-ray Adsorption Fine Structure (EXAFS), but both qualitative and quantitative data are still needed. In contrast, the reversible processes are easier to describe and are the subject of an important number of studies leading to data that is available and can be used for the prediction of retention processes. Moreover, most reversible processes, such as surface sorption processes, are kinetically faster and are generally the main processes controlling the concentration of the elements in the porewater, i.e. the elements which are mobile. In this work we will, therefore, focus on reversible surface retention processes, in which the surface processes are strictly related to a complex formation reaction between dissolved metal ions and the surface functional groups present at the surface of minerals.

How to describe the retention processes?

Adsorption, in which the ion of interest makes a chemical bond at the surface of the mineral, is one of the most relevant reversible processes responsible for the retention of radionuclides. There are two distinct models describing the equilibria of adsorption-desorption mechanisms at the mineral surface:

(1) empirical adsorption models, that are based on partitioning relationships with no predictive character; and (2) mechanistic adsorption models that describe the details of chemical species formation at mineral surfaces using thermodynamic formalisms; in this last case such models have predictive character but also require a better characterization of the system.

In the empirical model, the adsorption is often described by law of mass action and by simple distribution coefficients or Langmuir and Freundlich isotherms. The simplest distribution coefficients (KD) models (see example of Equation 1 in Chapter 2) describe the partitioning of solutes between the aqueous phase and the surface in a given set of experimental conditions. The KD is then constant according to the solid-to-liquid ratio in a given concentration range where the site saturation does not occur.

In more sophisticated empirical models, such as the Langmuir and Freundlich isotherms model, the sites of interaction are considered without specifying the type of species formed at the surface. Such models can explain the sites saturation (not possible with the KD model) and can therefore describe sorption isotherms in concentration. These models are widely used due to their conceptual and numerical simplicity, although they are seriously constrained as they are not able to describe chemical and physical perturbations in the geochemical system, such as the pH effect (Hochella and White,1990).

The mechanistic models are essentially extensions of the ion association models in aqueous chemistry that make the following assumptions: a) the surface functional groups react with dissolved solutes to form surface complexes; b) the physical-chemical equilibrium is described via mass action law equations with correlation factors applied for variable electrostatic conditions and boundaries; c) the surface charge is treated as a chemical reaction of the functional group and d) the apparent binding constants determined for the mass law equations are empirical parameters constrained by thermodynamic properties of the activity coefficients of the surface species (references in [START_REF] Davis | Surface Complexation Modeling in Aqueous Geochemistry[END_REF].

In this study, we will concentrate our attention only on mechanistic models focusing on three questions: What are the retention mechanisms to be considered? How to describe and quantify them and what are their limitations? How to interpolate the sorption mechanisms to environmental samples?

The retention mechanisms:

The main retention mechanisms occur at the surface of all mineral phases by the presence of surface hydroxide groups. They result from the hydroxylation of the surface upon exposure to water molecules and can be seen, by analogy with what happens in solution, as hydroxide surface sites. According to the type of mineral, they have different reactivities depending on the metal to which the OH group is bound, e.g. iron for hydroxides or silicium for silicates surfaces.

The adsorption of ions on the hydroxide surface sites is known as surface complexation. The ions can then interact with these sites leading to outer-(the ion keeps its first hydration sphere) or inner-(one or several water molecules in the first coordination sphere are replaced by atoms belonging to the reacting surface group) sphere complexes. For the clay minerals of interest in this work (e.g. 2:1 layer type structure), the hydroxide surface sites are located on the edges and can be identified as strong and weak sites. Such sites will be better described later.

In addition to the classical "surface complexation" mechanism, one has also to consider the ion exchange (Figure 30) that occurs in clay materials. The ion exchange comes from the permanent charge originated from isomorphous substitutions during the formation of the clay mineral. To compensate for the negative charge, cations are bonded by electrostatic interaction, and can be exchanged according to the composition of the solution. These cations are normally bounded at one type of sites denominated "basal sites". A schematic representation of the cation exchange sites existing at the basal and edge surfaces is given in Figure 30.

Besides the permanent negatively charged character, the charge of the mineral can either be positive or negative depending on the pH of the suspension due to the amphoteric nature of the surface functional groups. In acidic conditions, the edge surface becomes positive due to protonation of the functional groups, whereas in alkaline conditions the edge surface deprotonates and becomes negatively charged.

For clay materials, the combination of this pH-dependence charge with the permanent charge at the basal surface hinders the description of surface mechanisms, notably in terms of surface potential [START_REF] Ferrage | Influence of pH on the Interlayer Cationic Composition and Hydration State of Ca-Montmorillonite: Analytical Chemistry, Chemical Modelling and XRD Profile Modelling Study[END_REF][START_REF] Tournassat | Modeling Specific pH Dependent Sorption of Divalent Metals on Montmorillonite Surfaces. a Review of Pitfalls, Recent Achievements and Current Challenges[END_REF]. Description and quantification of surface mechanisms:

Cation exchange process

The cation exchange process is in most cases pH-independent (in normal environmental conditions) but it is strongly dependent on the ionic strength and chemical composition of the solution. Typically, a cation exchange reaction of metal ion B with valency z exchanging with a metal ion A with valency y on a clay mineral surface X is generically defined according to Reaction 18.

zA y (aq) + yB z Xz(s) ⇌ yB z (aq) + zAXy(s)

Reaction 18

Reaction 18 is defined by the selectivity coefficient K A/B X according to law mass action relationship (Equation 7).

𝐾 𝐴/𝐵 𝑋 = [AX(s)]z•γ B [B] y [B z X(s)]y•γ A [A] z Equation 7
In Equation 6, γA,B is the activity of the solute A and B (mol/L), the square brackets correspond to the concentration of A and B sorbed per unit mass divided by CEC (Cationic Exchange Capacity, eq/kg) or in solution (mol/L). For exchange metal ions, different conventions may be used as there is no unifying theory to calculate the activity coefficients [START_REF] Gaines | Adsorption Studies on Clay Minerals. II. A Formulation of the Thermodynamics of Exchange Adsorption[END_REF][START_REF] Gapon | On the Theory of Exchange Adsorption in Soils[END_REF][START_REF] Vanselov | The utilization of the base-exchange reaction for the determination of activity coefficients in mixed electrolytes[END_REF]. Amongst such conventions, the Gaines-Thomas convention is the most widely used.

Due to no unifying theory for the activity coefficient of the exchanger, the reference state of the activity coefficient of the adsorbed ions normally stands for 1, assuming that the exchanger is totally occupied by the same cation. This simplified assumption should be taken with caution since the selectivity coefficient can change according to the solution composition, as the activities coefficient are highly dependent on the ionic strength and the type of exchange sites occupied [START_REF] Mcbride | Interpretation of the Variability of Selectivity Coefficients for Exchange between Ions of Unequal Charge on Smectites[END_REF]. Therefore, the selectivity coefficient must be chosen carefully according to the system of interest [START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF]. Recently, there are models under development where the correction of activity for the different sites is considered (Reinoso-Maset andLy, 2014, 2016;[START_REF] Stammose | Sorption Mechanisms of Three Actinides on a Clayey Mineral[END_REF]. In this work, the Gaines-Thomas convention is used.

Surface complexation models and application to clay minerals

In contrast to the cationic exchange, the surface complexation is strongly dependent on the pH and weakly dependent on the ionic strength of the solution. In environmental conditions (pH approx. 7 and metal trace concentrations), surface complexation is the main process governing the retention of metals present in trace concentrations.

In the last decades different surface complexation models (SCM) have been proposed. The different models can be distinguished based on specific interfacial structures, resulting in various kinds of surface reactions and electrostatic correction factors to mass law equations. These models were originally developed to describe the surface reaction on oxides minerals, but they have been widely used to any mineral that develops an amphoteric surface charge, such as clay minerals [START_REF] Davis | Surface Complexation Modeling in Aqueous Geochemistry[END_REF]. In SCM models, as for the empirical models, the parameters must be determined in a wide range of solution compositions in order to reproduce the natural conditions as closely as possible [START_REF] Parks | Surface Energy and Adsorption at Mineral-Water Interfaces: An Introduction[END_REF].

The SCM models can be described as chemical equilibrium approaches defined by surface species, chemical reactions, equilibrium constants, mass balances and charge balances [START_REF] Goldberg | Surface Complexation Modeling[END_REF]. Different models have been defined in the literature based on the reactive surface function group, that undergoes both protonation and deprotonation reactions, with common characteristics and adjustable parameters. In acidic conditions, there is protonation of the functional groups (Reaction 19), with a release of H + ions to the solution changing the pH of the media. In more alkaline conditions the deprotonation of the functional groups is favoured (Reaction 20). The equilibrium constants of the protonation and deprotonation reactions are shown, respectively, in Equations 7 and 8 below,

K + ≡SOH = ( [≡SOH 2 + ] [≡SOH]•{H + } ) • 𝑒 (FΨ/RT) Equation 8 K - ≡SOH = ( [≡SO -]•{H + } [≡SOH] ) • 𝑒 -(FΨ/RT) Equation 9
In Equation 8 and Equation 9, the F is the Faraday constant (96485 C•mol -1 ), Ψ is the surface potential, R the molar gas constants, and T the absolute temperature. The squared brackets correspond to the concentration of the respective surface functional group and the curly brackets correspond to the aqueous activities. The aqueous activities are calculated as a function of the ionic strength of the solution according to the Extended Debye-Hückel, Davies, SIT or Pitzer models [START_REF] Grenthe | Chapter IX Estimations of Medium Effects on Behaviour Of Uranium in Callovo-Oxfordian Formation At 20 ºC Flávia Maia | IMT Atlantique Nantes Thermodynamic Data[END_REF]).

The best known for this type of models is the Davies equation, which is suitable for ionic strength below 0.1 M and will be the one considered hereafter.

Moreover, different models were proposed to describe the surface potential. These models differ on the structural representation of the solution-surface interface: CCM─ Constant Capacitance Model [START_REF] Schindler | Die aciditat von silanolgruppen[END_REF][START_REF] Hohl | Interaction of Pb 2+ with hydrous α-Al2O3[END_REF][START_REF] Sposito | The Surface Chemistry of Soils[END_REF][START_REF] Schindler | The Surface Chemistry of Oxides, Hydroxides, and Oxide Minerals[END_REF]; DLM─ Diffuse Layer Model [START_REF] Stumm | Specific chemical interactions affecting the stability of dispersed systems[END_REF][START_REF] Huang | Specific adsorption of cations on hydrous α-Al2O3[END_REF] and TLM─ Triple Layer Model [START_REF] Hayes | In-situ Xray absorption study of surface complexes: Selenium oxyanions on α-FeOOH[END_REF]. When the surface functional group is characterised by either one or two protons, the SCM model is developed from the Stern Model based on one layer at surface, so-called Stern layer. The expansion of the TLM models is known as Charge Distribution Multi-Site Complexation (CD-MUSIC), which considers charge distribution bond valence principles and identifies different types of surface sites based on their crystallographic structure. The location and surface configuration of the adsorbed ions for the different SCMs models are schematically shown in Figure 31. In the case of clay minerals, specially 2:1 type clay minerals, even though SCM model approaches can describe sorption experimental data the thermodynamic parameters obtained have no real significance as the correction of the electrostatic potential is not correct as the potential arises from two types of charges that are not located structurally in the same place (i.e. at the basal and edge surfaces).

For this reason, specific models are being developed for surface complexation on clay minerals, and recently a model considering the spill-over effect on the surface potential was applied successfully for describing the retention of U(VI) on montmorillonite (Tournassat et al., 2018). However, this research is still in progress and this type of model cannot yet be used in this work due to the limitation of thermodynamic data available for the other clay minerals.

An alternative approach is to ignore the electrostatic parameter and use Non-electrostatic surface complexation models (references in [START_REF] Parks | Surface Energy and Adsorption at Mineral-Water Interfaces: An Introduction[END_REF]Bradbury andBaeyens, 1997, 2005c;Marques Fernandes et al., 2008, 2012;[START_REF] Poinssot | Experimental and modelling studies of caesium sorption on illite[END_REF][START_REF] Turner | Surface-charge properties and UO2 2+ adsorption of a subsurface smectite[END_REF]. In clay mineral systems, nonelectrostatic models were originally developed for Zn/Ni-montmorillonite systems as the electrostatic term was not required for describing titration experiments. Such models have been widely used by Bradbury and Baeyens (1997, 2000[START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF][START_REF] Donat | The Removal of Uranium (VI) from Aqueous Solutions onto Natural Sepiolite[END_REF] for different metal cations and radionuclides in clay systems. In 1997, Bradbury and Baeyens proposed a non-electrostatic model approach and they defined two protolysis sites (protonation/deprotonation sites) to describe the surface reactivity as a function of the pH for 2:1 clay minerals. Besides the protolysis sites, an additional strong surface site of low capacity had also to be defined to describe the surface complexation of metal ions. Moreover, cation-exchange processes were also required to describe the sorption behaviour mainly at low pH, high equilibrium metal concentration or low ionic strength. This model approach is called the 2 Sites Protolysis Non-Electrostatic Surface Complexation and Cation Exchange (2SPNE SC/CE) sorption model. Since the development of this model, numerous parameters have been obtained for several elements (e.g. Eu, Th, U). In 2005 and 2009, Bradbury and Baeyens demonstrated a linear free energy relationship (LFER model) between the equilibrium constant of the surface complexation and the hydrolysis stability constants for different cations for the strong sites of the surface of montmorillonite (Bradbury and Baeyens, 2005c) and illite (Bradbury and Baeyens, 2009a). The combination of the non-electrostatic surface models and the LFER allows the prediction of all cations´ adsorption on most clay minerals considering the competitive sorption of co-existing cations as well (Bradbury and Baeyens, 2011;Marques Fernandes et al., 2012).

In this work, we will use the 2SPNE SC/CE model to describe the sorption behaviour of U(VI).

How to interpolate sorption models on natural samples?

The quantification of adsorption processes in natural systems can be done through two approaches: the component additivity (CA) and the generalised composite (GC) approach [START_REF] Davis | Application of the Surface Complexation Concept to Complex Mineral Assemblages[END_REF][START_REF] Waite | Approaches to Modelling U(VI) Adsorption on Natural Mineral Assemblages[END_REF].

The CA approach, known also as bottom-up approach, attempts to model the sorption processes taking into account all minerals constituting the natural system and assumes that each mineral has unique surface sites with characteristic sorption properties. In this case, the characterisation of the surface of each single mineral is necessary. For practical reasons, this approach requires the simplification of the system in order to consider only the most important interacting phases. However, in some cases, the CA approach can lead to an over simplification of the systems, where the sorption is controlled only by one mineral-phase (NEA, 2005;[START_REF] Payne | Guidelines for Thermodynamic Sorption Modelling in the Context of Radioactive Waste Disposal[END_REF].

In the semi-empirical model (GC approach), it is assumed that the mechanistic description of surface composition of the mineral assemblage is inherently too complex to evaluate the contributions of individual phases. Instead, generic surface functional groups are assigned to the mineral mixture, in which the stoichiometry and formation constants for each sorption equilibrium and site capacities are determined by fitting experimental data. The GC approach is merely a fitting rather than a predicting procedure [START_REF] Davis | Application of the Surface Complexation Concept to Complex Mineral Assemblages[END_REF]Marques Fernandes et al., 2015;[START_REF] Waite | Approaches to Modelling U(VI) Adsorption on Natural Mineral Assemblages[END_REF]. These kind of model approaches have been widely used to predict the behaviour of different cations in different mineral-phases (e.g. Marques Fernandes et al., 2015;Chen et al., 2014).

For clay rock systems, bottom-up approach has been applied considering the 2SPNE SC/CE model. The bottom-up approach was applied successfully on Opalinus clay (Bradbury et al., 2011;[START_REF] Hartmann | Sorption of Radionuclides onto Natural Clay Rocks[END_REF][START_REF] Van Loon | The Sorption Behaviour of Caesium on Opalinus Clay: A Comparison between Intact and Crushed Material[END_REF] for several metals (Ni(II), Co(II), Eu(IIII), Th(IV) and U(VI)). It was also possible to describe Cs/COx clay (Chen et al., 2014) and Ni/COx clay (Chen et al., 2014a) considering that the clay fraction was characterised as illite and interstratified I/S, where I/S is mixture of illite and montmorillonite as proposed in [START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF]. For the latter case, Chen et al.

(2014a) observed some limitations in the interpolation for natural systems (COx claystone) highlighting the difficulty of generalising the quantitative description of mineral reactivity.

At this stage, and considering the objective of this work, we propose to use the 2SPNE SC/CE model with the bottom-up approach. The simplifications and assumptions correspond to those already published [START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF]Bradbury et al., 2011;Chen et al., 2014Chen et al., , 2014a;;Marques Fernandes et al., 2015[START_REF] Jin | Surface Complexation Modeling of U(VI) Adsorption on Granite at Ambient/elevated Temperature: Experimental and XPS Study[END_REF] and are as follows:

i. 2:1 clay minerals control the retention; this corresponds to a content of 30-40% in COx formation (mainly composed by illite and interstratified I/S layers). ii. I/S mixed layers should be treated as sorption phase separately, although due to a lack of data and sorption models, here, the I/S interstratified layers will be assumed as two phases, where each phase corresponds to 50 %wt of illite and smectite, respectively. iii. Montmorillonite will be used as a proxy for smectite phases since it is a well-known and studied smectite. iv. The potential contributions of chlorite and kaolinite will not be taken into account on the overall sorption of uranium.

The sorption parameters for both illite and montmorillonite are summarised in Table 18, and the surface complexation reactions and their respective selective coefficients are provided in Table 19.

For both clay minerals, the number of surface and exchange sites and their respective mass are detailed in Appendix B.2. The thermodynamic data of the aqueous speciation and solids are included in the Thermochimie database and a summary of the most relevant data is described in Appendix E.

Concerning the aqueous Ca-U(VI)-CO3 ternary complexes, the new thermodynamic data presented in the previous chapter will be used in the geochemical model (Table 14 in Chapter 2). In COx systems, Mg-U(VI)-CO3 ternary complexes [START_REF] Lee | Formation, Stability and Structural Characterization of Ternary MgUO2(CO3)3 2and Mg2UO2(CO3)3(aq) Complexes[END_REF] were also taken into account in the model. Based on the 2SPNE SC/CE model (Table 18 andTable 19) and taking the COx formation chemical conditions into account (pH = 7.2 and pCO2 = 10 -2.1 atm, see Table 2 in Chapter 1), it is interesting to notice that the surface complexation is controlled mainly by the sorption of carbonato complexes of U(VI) on montmorillonite (Table 20) while the corresponding constants for illite are unknown. The importance of the carbonate complexes on the sorption process was observed on different model phases from hydro(oxides) to different clay minerals and clay rocks These modelling studies predict that U(VI) adsorption decreases at alkaline pH when carbonate anions are present due to the formation of strong aqueous uranyl-carbonate complexes [START_REF] Davis | Approaches to Surface Complexation Modeling of Uranium(VI) Adsorption on Aquifer Sediments[END_REF]. However, the impact of the aqueous carbonate complexes on U(VI) sorption depends on whether these complexes adsorb on the surface of clay minerals. In Marques Fernandes et al. (2012), the authors found necessary to include ternary uranyl-carbonate-surface (U(VI)-CO3-S) complexation reactions to describe their experimental data on the sorption of U(VI) onto montmorillonite, as a function of different carbonate contents. The formation of such ternary surface complexes was also suggested by [START_REF] Waite | Uranium(VI) Adsorption to Ferrihydrite: Application of a Surface Complexation Model[END_REF] for describing the sorption of U(VI) on ferrihydrite. However, the existence of ternary U(VI)-CO3 surface complexes has been questioned due to a lack of strong spectroscopic evidences (Tournassat et al., 2018). Tournassat and co-workers (2018) have been developing a model that integrates the electrostatic correction and have recently suggested that such surface complexes are not necessary to describe U(VI)/montmorillonite systems.

Most studies have focussed on U(VI)/montmorillonite systems whereas only a few works have been reported on the sorption mechanisms of U(VI) on illite (Bradbury et al., 2005c;Gao et al., 2015;Marques Fernandes et al., 2015) and to our knowledge no studies have been published concerning the effect of carbonates on the sorption of U(VI) on illite. Therefore, if carbonato surface complexation on illite surface exists, the only parameter missing in the model is the stability constants of the U(VI)-CO3 surface complexation reactions, and in case these are needed they will be fitted according to the experimental data.

Experimental work

The objective of the present work is to apply the bottom-up approach proposed by Bradbury andBaeyens (1997, 2011) on Callovo-Oxfordian formation by following a stepwise approach:

-The first step is to generate experimental data for illite in order to test the model and fit, if necessary, any surface parameters, particularly in the presence of carbonates. -The second step is to test the model on the COx clay fraction, where illite and I/S are the main phases responsible for the retention of radionuclides on COx formation. -The third and final step is to test the model on a "real system" using the COx claystone natural samples.

The parameters of interest for solution characterisation are the pH, alkalinity and the initial concentration of calcium. The experimental cases defined in the present study are described in Table 21. As proposed in Chapter 2, for all modelling exercises both the pH and the alkalinity were fixed to those experimentally measured. The pCO2 was let free accordingly to the pH/alkalinity measured. The illite used for the U(VI) sorption studies is known as Illite du Puy from the Oligocene geological Formation in Le Puy-en-Velay (Haute-Loire, France). The purification and conversion into the homoionic Na form was performed and provided by BRGM (Bureau de Recherches Géologiques et Minières, France).

The claystone sample was collected and provided by Andra in 2008 and corresponds to the borehole sample EST51769 of Callovo-Oxfordian claystone natural sample. The sample was stored in an inert atmosphere of N2(g) to avoid oxidation and carbonatation. The claystone is composed of 17 wt% of illite, 28 wt% of I/S, 2 wt% of kaolinite, 3 wt% of chlorite, 25 wt% of calcite and dolomite, 25 wt% of quartz and feldspar and 0.8 wt% of pyrite (personal communication J. C. Robinet).

The Callovo-Oxfordian clay fraction corresponds to the fraction sized below 2µm of the Callovo-Oxfordian claystone (Appendix A.6). The sample was purified and provided by BRGM and is composed by of 37.8 wt% illite and 62.2 wt% of interstratified illite/smectite.

Prior to the sorption experiments, we performed kinetic adsorption tests on COx claystone natural samples and illite to evaluate the time needed to reach steady-state equilibrium (details of the experimental procedure can be found in Appendix A.7.2). The sorption of U(VI) on COx natural sample (Figure 32a) and illite (Figure 32b) was evaluated for 14 days. As can be observed in Figure 6, the sorption of uranium in the COx claystone natural sample in synthetic COx porewater is much lower compared to the illite in ionic media (0.1 M NaCl). For the experiments of sorption of U(VI) under the synthetic Callovo-Oxfordian porewater conditions, after 3 days of contact, we assumed that the steady-state was reached. In the case of illite, the equilibrium appeared to be established after one day of contact. To ensure that steady-state is always reached independently of the chemical conditions, 3 days of contact were considered for both illite and Callovo-Oxfordian claystone sample, as also proposed for other sorption studies (Bradbury et al., 2005a;Bradbury et al., 2009a;[START_REF] Chisholm-Brause | Speciation of Uranyl Sorbed at Multiple Binding Sites on Montmorillonite[END_REF]Joseph, 2013;Tournassat et al., 2018).

The experimental conditions for both sorption edge and sorption isotherms are described in Table 22. )]initial (M) 10 -7 10 -7 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5

Uranium sorption edge and isotherm experiments

We measured the sorption isotherms in concentration for illite, COx clay fraction and COx claystone natural samples in equilibrium with inert atmosphere-Ar(g), atmospheric pCO2, 1% of CO2(g) and 5% of CO2(g). The pH of the synthetic solutions in equilibrium at the different atmospheres and alkalinity were measured at the end of the experiments. In inert atmosphere, considering the strong influence of carbonate on U(VI) speciation, alkalinity was measured even considering the a priori inert conditions. The samples were prepared at solid-liquid ratio of 2, 10 and 40 g/L (clay fraction, illite and claystone, respectively) in 0.1 M of NaNO3 or synthetic COx porewater. For each experiment, a series of blank samples were used to control for uranium sorption/precipitation in tubes and filters, and to control for the alkalinity of the system. The details of the experimental procedure and the material used are described in Appendix A and the experimental results are detailed in Appendix C.5.1.

We carried out sorption edge measurements (initial concentration of U(VI)total ≈ 10 -7 M) as a function of pH in equilibrium with Ar(g) and atmospheric pCO2 to quantify the influence of carbonates on the sorption of uranyl onto illite. In inert atmosphere, the pH of the system was controlled by adding 0.01 M of HNO3 or free carbonated NaOH directly to each sample. For the experiments in atmospheric conditions, the pH of each stock solution was equilibrated in atmospheric conditions by overnight agitation (more details of the procedure are provided in Appendix A.7) until the pH of interest was reached. In these experiments, the carbonate content was determined by alkalinity titration measurements. The samples were prepared at solid-liquid ratio of 0.5-1 g/L (depending of the pH of the system) in 0.1 M of NaNO3 (inert atmosphere) or NaCl (atmospheric pCO2). For each experiment, a series of blank samples were used to control for uranium sorption/precipitation in tubes and filters and to control for the alkalinity of the system. The details of the experimental procedure and the material used are described in Appendix A and the experimental results are detailed in Appendix C.5.2.

Sorption of U(VI) on illite

In the absence of carbonate

The sorption of U(VI) takes place mainly on the reactive sites associated to the surface hydroxyl groups (≡SOH) located along the edges of the clay platelets, in which it can protonate and deprotonate, as a function of the pH, and form surface complexes with aqueous uranyl species (Marques Fernandes et al., 2015). In the absence of carbonates, the distribution of the aqueous speciation of the uranyl (Figure 33a) as a function of pH in 0.1 M of NaNO3 is controlled only by hydrolysis species, which are the most relevant for the sorption of uranyl (Figure 33b)

As expected, the sorption of uranyl is clearly dependent on the pH of the solution; the KD increases with increasing pH , reaching a higher KD of U(VI) at neutral pH conditions (Figure 34). As can be seen in Figure 34, the KD obtained (open circles) is similar to the one reported by Bradbury et al., 2009b (black squares). Despite the KD trend being the same and in the same order of magnitude, the KD obtained in the present work is slightly lower for pH below 6, whereas slightly higher at pH range 6-8. Above pH 8, the results are similar. In the present work, using the surface complexation reactions of hydrolysis species in both strong and weak sites (Table 19), the sorption model (black solid line) reproduces relatively well the experimental curve increase. When comparing the sorption isotherms of the present work (black circles, Figure 36) with the data reported by Bradbury and Baeyens (2005) (blue open squares, Figure 36), and despite the different pH at equilibrium, both experimental data (dots) and model approaches (solid lines) are in agreement.

The model approach of this work (black solid line) not only reproduces the present experimental data but also the experimental data reported by Bradbury and Baeyens (2005).

In the presence of carbonate

In the presence of carbonates, the sorption of U(VI) onto illite is clearly affected at pH > 6 in comparison with the data obtained in the absence of carbonates (Figure 37). The titration measurements of the alkalinity in the stock solutions and pH measured before and after the equilibrium with illite are shown further in Table 23.

The sorption edge experiments in the presence of carbonates in equilibrium with pCO2 atmospheric (blue circles) indicate that the maximum KD of U(VI) is reached at pH 6.5, one order of magnitude lower to the KD observed in the absence of carbonates (open circles, Figure 37). Furthermore, whereas the KD is nearly constant at pH > 7 in the absence of carbonates a strong decrease is observed in the presence of carbonates.

In the present work, the model approach considers: a) the surface reactions for illite described in Table 19, where the carbonato surface complexation reactions are not considered (due to lack of data); b) the alkalinity measurements performed in the beginning of the experiments; and c) the pH measurements of the equilibrium. Here, it was assumed that alkalinity does not change significantly after the equilibrium, i.e. the alkalinity is constant. As shown in Figure 38, the presence of carbonates strongly affects the speciation of uranyl in solution. This effect is more pronounced at high pH, as shown in Figure 38a while at pH 6, the speciation of U(VI) is controlled by uranium-carbonate aqueous species. In conclusion, we can assume that there is a competition for U(VI) between its complexation in surface and in solution (Figure 38b).

i Calculations in Phreeqc from the equilibrium between pH measured at the end of the experiments; maximum and minimum considering the uncertainty of the alkalinity measurements The modelling results indicate that it is possible to reproduce relatively well the experimental sorption results of U(VI) at neutral pH (see further Figure 39a). In more acidic conditions (pH ≈ 5), as already observed in the absence of carbonates, the model slightly overestimates the sorption of U(VI) on illite.

On the other hand, in more alkaline conditions (pH ≈ 8.1), the model underestimates the sorption behaviour of U(VI).

We can observe by evaluating the pCO2, for these experiments (calculated based on the relation pH/alkalinity), that it is not always as expected (pCO2 atmospheric = 10 -3.5 atm). As shown in Table 23, the calculated pCO2 is slightly higher at pH < 7 than pCO2 atmospheric, affecting the calculations. However, the model that considers equilibrium with atmospheric pCO2 can reproduce well the data, indicating that we are not so far from the "theoretical" water/gas equilibrium.

Since it is not possible to reproduce the experimental data at pH > 7, these results raise the question of whether to use or not the uranyl-carbonato surface complexation reactions. Since there is no thermodynamic data relative to the uranyl-carbonato surface complexations reaction onto illite, as a first approach, we considered the same data as for montmorillonite (see Table 19), assuming that both minerals have similar reactivity.

The results of the model approach are shown in Figure 39b. The agreement between the model and the experimental data below pH 7 is relatively good. When the pH is above 7 and the uranyl carbonate species are the major aqueous species in solution, even considering the thermodynamic constants in the U(VI)-CO3 surface reactions as the ones for montmorillonite, the model still underestimates the sorption of U(VI) onto illite (Model I).

These results led us to question the reliability of the values of the surface reaction constants for uranyl carbonates proposed for montmorillonite on the surface edges of illite. Thus, we proposed a second approach by changing the stability constants of the uranyl carbonate surface reactions. The number of variables implicated in this approach must be restricted to avoid high uncertainty. For this reason, only one stability constant has been selected, i.e. that associated to the predominant complex in the experimental conditions of this work, ≡S s OUO 2 (CO 3 ) 2 3-. The fitting of the stability constants was done manually, and its best fit was found to be log10 K = 17.5 ± 0.5 (Figure 39c). In the following calculations, the uncertainty of this constant will be considered in the prediction approach.

The ≡S s,w OUO 2 CO 3 -reactions were also evaluated on their ability to influence the sorption behaviour of U(VI) in the presence of carbonates. For that, Log10K of 𝑆 𝑠 𝑂𝑈𝑂 2 𝐶𝑂 3 -and S w OUO 2 CO 3 -was defined from the montmorillonite as 9.8 and 9.3, respectively. In Figure 39d, we show the impact of considering or not these two additional reactions in the model approach. In Model 0 neither ≡S s OUO 2 (CO 3 ) 2 3-nor ≡S s,w OUO 2 CO 3 -are considered; in Model II only ≡S s OUO 2 (CO 3 ) 2 3-with log10 K = 17.5 (±0.5) is considered and in Model III the ≡S s OUO 2 (CO 3 ) 2 3-, the ≡S s OUO 2 CO 3 -and ≡S w OUO 2 CO 3 -are considered, with log10 K of 17.5, 9.8 and 9.3 respectively. We can observe that the presence of ≡S s,w OUO 2 CO 3 -does not affect the sorption of U(VI) on illite as a function of pH, indicating that ≡S s OUO 2 (CO 3 ) 2 3-plays the major role on the sorption of U(VI) at high pH.

Next, we evaluated the reliability of the stability constants of ≡S s OUO 2 (CO 3 ) 2 3-based on sorption isotherms of U(VI) at different ranges of pCO2. Despite the fact that ≡S s,w OUO 2 CO 3 -does not have a role on the sorption behaviour of U(VI) on illite, as a function of pH in equilibrium with atmospheric pCO2, it may have an impact on the sorption of U(VI) at higher concentrations of carbonate in solution. In this case, it will also be evaluated whether it should or not be included in the model. The model approach will follow Model III as shown in Figure 39d and Table 24. Sorption isotherms experiments performed in presence of different concentration of carbonates (atmospheric conditions and 1% and 5% of CO2(g)) were used to evaluate the influence of the ≡S s,w OUO 2 CO 3 -complexes. The results of the sorption isotherms of U(VI) in equilibrium with pCO2 atmospheric are shown in Figure 40. The sorption of U(VI) onto illite does not seem significantly affected when the carbonates are added to the system at < 10 -8 M of U(VI) in equilibrium in a solution with pH 7.3. However, the experiments performed in the absence of carbonates are not completely free of carbonates. Besides the low contamination observed in the latest experiments, the concentration of carbonates (7•10 -5 eq/L) is relatively close to the one measured in atmospheric pCO2 (alkalinity of 2•10 -4 eq/L of carbonates). The same issue was also observed on the experiments performed on montmorillonite reported by Tournassat et al., (2018). The model approach, using the stability constants of the carbonato surface complexation reactions proposed in this work for illite (see Table 24), reproduces well the experimental sorption of U(VI) in the presence of carbonates (atmospheric pCO2) (Figure 41b), whereas considering only one uranyl carbonate surface complexation (≡S s OUO 2 (CO 3 ) 2 3-) to control the sorption of U(VI)-CO3 complexes slightly underestimates the sorption of U(VI) (Figure 41a). Both model approaches, considering the uncertainty of the stability constants, can reproduce the behaviour of U(VI) in illite in equilibrium with pCO2 atmospheric. However, it is difficult to define a selection criteria for the best model approach to describe the behaviour of U(VI) in the presence of low concentration of carbonates.

Finally, we performed two complementary sorption isotherms at partial pressure of 1% and 5% of CO2(g). The final pH and alkalinity for each experiment are shown in Table 25. The results show that in the presence of higher concentrations of carbonates in solution the sorption of U(VI) on illite slightly decreases (comparison between air and 1%-5% of CO2(g)). However, the behaviour of uranyl does not seem to be affected significantly by changing the conditions from 1 to 5% of CO2(g), even when the pH decreases (Figure 42). The pH decrease observed does not affect the U(VI) sorption, which means that the U(VI) behaviour is not only controlled by the U(VI)-CO3 aqueous species. These results clarify the importance of considering also the sorption mechanisms of the U-CO3 ternary surface complexes on illite. The prediction model was used to quantitatively describe U(VI) sorption on illite in the presence of carbonates in equilibrium with 1% and 5% of CO2(g) (Figure 43). As shown previously, Model II considers only one carbonate surface complexation reaction (Figure 43a,b), whereas Model III considers the three reactions (as for montmorillonite) (Figure 43c,d). Both models take into account the fitted stability constant for the ≡S s OUO 2 (CO 3 ) 2 3-and its respective uncertainties.

The results from the model indicate that both approaches might be suitable to explain the sorption of U(VI) in the presence of higher concentration of carbonates in solution. Nevertheless, if we look in more detail Model II might be a better approach than Model III, in which all constants for carbonato surface complexes are considered. Model II reproduces better the sorption of U(VI) at higher concentrations of U(VI) as shown when seeing the results in 5% of CO2(g) (Figure 43c andd). Considering this, different experiments (e.g. changing the range of concentration of carbonates in the system) should be done to ensure the reliability of such surface species in the fitting of the parameters when in presence of a large experimental dataset. Alternatively, one could perform surface spectroscopic analysis to validate the existence of the U(VI)-CO3 surface complexes. However, this is not the subject of the present work. In order to simplify and decrease the number of uncertainties, Model II will be chosen for the following work, in which only one carbonate surface complexation (≡S s OUO 2 (CO 3 ) 2 3-) will be considered in the illite phase model. 24. Error bars that are not visible are smaller than the symbol size.

Impact of the presence of Ca-U(VI)-CO 3

We next studied the impact of the formation of calcium-uranyl-carbonate ternary complexes on the sorption of U(VI) onto illite. The experimental conditions were set as closely as possible to the ones that characterize Callovo-Oxfordian formation at 20ºC in equilibrium with 1% of CO2(g) (see synthetic COx porewater composition in Table C. 7). Due to the possible decrease of KD of U(VI) onto illite, expected from the formation of stable Ca-U(VI)-CO3 species in solution, the solid-liquid ratio was adjusted to 10 g/L to better quantify the concentration of adsorbed U(VI). The pH and alkalinity at equilibrium were measured at 7.1 (± 0.02) and 2.5 (± 0.1) meq/L, respectively. The results of the sorption isotherm of U(VI) indeed show that the formation of Ca-U(VI)-CO3 complexes does not favour the sorption of U(VI) (Figure 44) as compared to when we performed the experiment at 1% of CO2 (g) in the absence of Ca. The results of the model considering the U(VI)-CO3 surface complexation reactions given in Table 24 and the parameters of the Ca-U(VI)-CO3 complexes proposed in the present work (Table 14 in Chapter 2 In summary, the presence of carbonates in solution influences the sorption of U(VI) onto illite, in which a sorption decrease is observed due to the formation of U(VI)-CO3 aqueous species. However, the surface complexation reactions for U(VI)-CO3 complexes must be included in the model to quantitatively explain the experimental data. The stability constants of the main surface complex (≡S s OUO 2 (CO 3 ) 2 3-) were therefore fitted from the sorption edge experimental data, in which we obtained the proposed value of log10 K = 17.5 ± 0.5. With this new stability constant, it was then possible to predict the behaviour of U(VI) at pH ≈ 7 for a wide range of pCO2 (atmospheric conditions, 1% and 5% of CO2(g)) and in the "COx conditions" (1% of CO2(g) in the presence of calcium in solution).

The next step is to assess whether Model II, together with the one proposed for montmorillonite (Table 3), can predict the sorption data on the COx clay fraction.

Sorption of U(VI) on Callovo-Oxfordian clay fraction

The clay fraction of Callovo-Oxfordian claystone is composed mainly by illite and interstratified I/S. Kaolinite or chloride may be present but will be considered negligible. As a reminder, in the present study both illite and montmorillonite are considered as model phases for the modelling and I/S corresponds to a mixture of illite and smectite. In this study we evaluated the retention of U(VI) in the absence of carbonates, in the presence of carbonates in equilibrium with 1% of CO2, and under Callovo-Oxfordian conditions, i.e. in the presence of the Ca-U(VI)-CO3 ternary complexes in solution.

We only measured the sorption isotherms in concentration at the pH values relevant to the Callovo-Oxfordian formation.

In the absence of carbonate

The sorption isotherm of U(VI) for COx clay fraction in the absence of carbonates is shown further in Figure 46a. The pH of equilibrium was 7.0 (± 0.2) and the alkalinity was found to be 0.2 (± 0.09) meq/L, almost 70 % higher than that measured in the illite system. From Figure 46a we can also compare the sorption of U(VI) on illite and clay fraction in the absence of carbonates. The trends are similar, although with a weaker retention for COx clay fraction than illite in the conditions studied, in concentrations of U(VI) below 10 -8 M.

Independently of the possible contamination of carbonates, the model cannot completely reproduce the data as it overestimates the retention for the lowest concentrations of U(VI). However, there is a relatively good agreement between the calculation and the experimental data for the highest concentrations. The sorption experiment using clay fraction in the absence of carbonates was performed only once, and for this reason we do not have any possible explanation for such difference. Nevertheless, we decided to evaluate the model in the most realistic conditions possible, which, in this particular case, meant that the U(VI) behaviour is also governed by the presence of carbonates.

In the presence of carbonate

We next measured the sorption isotherms in concentration on clay fraction in the presence of carbonates in equilibrium with 1% of CO2(g). The pH of equilibrium was 7.2 (± 0.01) and the alkalinity was 2.4 (± 0.9) meq/L. Comparing the results obtained for illite and clay fraction, no significant differences were observed between these two phases (Figure 46b).

The model approach, using the parameters given in Table 19 and Table 24, can reproduce well the U(VI) sorption on the clay fraction (Figure 46b red dash-line). The model results indicate that considering illite and montmorillonite to describe the interstratified I/S is a reliable approach.

Impact of the presence of Ca-U(VI)-CO 3

We performed the sorption experiments of U(VI) onto COx clay fraction in equilibrium with synthetic COx porewater that we previously prepared (see Appendix A.5 and Table C. 7.). In the presence of carbonates and calcium in the system, the sorption isotherm of U(VI) was measured at a pH of 7.1 (± 0.02) and alkalinity of 2.4 (± 0.9) meq/L. In Figure 47a, we can compare these results with those measured in the absence of Ca in equilibrium with 1% of CO2(g). The formation of Ca-U(VI)-CO3 promotes, as observed also in the illite system, a higher decrease of U(VI) sorption when comparing the experimental data where only U(VI)-CO3 complexes are formed. As observed in the absence of calcium in the system, the sorption behaviour of U(VI) does not change comparatively either in illite or in the clay fraction (illite + I/S) (symbols in Figure 47b).

The proposed model approach, considering the surface parameters both on illite (Model II) and montmorillonite can reproduce well the behaviour of U(VI) in Callovo-Oxfordian conditions (green dashed-lines Figure 47b).

So far, we have shown that the model works in conditions relevant to the Callovo-Oxfordian conditions, i.e. in the presence of carbonates and Ca at pH ≈ 7. We have also shown that using illite and montmorillonite as model phases to assume the I/S interstratified layers are usefull to model the COx clay fraction. However, the agreement is not perfect since the model could not predict well the experiments performed in "inert" conditions. Considering the objectives of this work, we decided not to focus on this point and to follow the objective of understanding the U(VI) behaviour in natural systems, i.e. closest to real conditions, and assess the reliablity of the third important assumption, which is interpolating the model used to the Callovo-Oxfordian claystone natural samples. Flávia Maia | IMT Atlantique Nantes

Sorption of U(VI) on Callovo-Oxfordian claystone

The COx claystone (as natural sample), as mentioned previously, is composed mainly of clay minerals (illite and interstratified I/S layers), carbonates (calcite) and quartz. The sorption isotherm of U(VI) on natural samples was performed in equilibrium with a synthetic COx porewater. This porewater was prepared as closely as possible to the composition of the in situ groundwater in the presence of 1% CO2(g) (see Appendix A.5 and Table C. 7). Due to the possible decrease of U(VI) distribution coefficient on natural samples, expected from the formation of Ca-U(VI)-CO3 in solution, the solid-liquid ratio was adjusted to 40 g/L to better quantify the concentration of U(VI) adsorbed. The pH of equilibrium was 7.5 ± 0.1 and the alkalinity was 3.3 ± 0.6 meq/L.

The results of the sorption isotherm of U(VI) in COx natural samples are shown in Figure 48 (whitesquares). In comparison with the previous results (illite and clay fraction, see above Figure 47b), there is a significant decrease of U(VI) sorption when natural samples of COx formation are considered. This is coherent with the assumption that only the clay fraction is responsible for the retention, since it represents 45 wt% of the COx natural sample and Ca-U(VI)-CO3 aqueous complexes are present.

The results of the model approach considering the surface reactions for illite and montmorillonite and the thermodynamic data for Ca-U(VI)-CO3 provided in Table 14 (see Chapter 2) are shown in Figure 49. The model approach reproduces well the experimental data in Callovo-Oxfordian conditions which lie within the uncertainties of the prediction model associated to the surface complexation reaction of ≡S s OUO 2 (CO 3 ) 2 3-species. According to the model calculations, the behaviour of U(VI) is controlled mainly by its aqueous speciation (Table 26). The surface complexes also play an important role on the behaviour of U(VI), in that the U(VI)-CO3 surface complexes on the surface of illite are responsible for the sorption of U(VI) in Callovo-Oxfordian formation. The impact of Ca-U(VI)-CO3 ternary complexes on the sorption of U(VI) on clayey rocks was already reported by [START_REF] Hartmann | Sorption of Radionuclides onto Natural Clay Rocks[END_REF] on Callovo-Oxfordian, Joseph et al. (2011) and Amayri et al. (2016) on Opalinus clay, and Marques Fernandes et al. (2015) on Opalinus and Boda clay.

For [START_REF] Hartmann | Sorption of Radionuclides onto Natural Clay Rocks[END_REF], the sorption of uranyl on COx claystone was studied in 0.1 M of NaClO4 in equilibrium with atmospheric pCO2 as a function of pH. In the model approach proposed by Hartmann and co-workers, the U(VI)-CO3 surface complexation reactions were not included, and this could explain why they were not able to predict the sorption of uranyl at pH between 6 and 9 (Figure 50, black line model). When those surface complexation reactions are considered, the model improves considerably (Figure 50, red lines model). Note that this improvement of the model is not only linked to the U(VI)-CO3 surface reaction considered for illite but also to new the thermodynamic data of Ca-U(VI)-CO3 complexes (see Table 14 in Chapter 2) (they considered the stability constants for Ca-U(VI)-CO3 complexes reported by Bernhard et al., 2001). It should also be noted that no measurements of carbonates were done by Hartmann and co-workers, so for modelling purposes, the concentration of carbonates was controlled by pCO2 of 10 -3.5 and calcite. When comparing the experimental results of Opalinus and Boda clay with Callovo-Oxfordian clays (present work), the sorption of U(VI) is lower for the Callovo-Oxfordian natural samples (Figure 51). Here, we can assume that the synthetic porewater composition plays a major role in the sorption mechanisms of U(VI). The main parameters that distinguish the synthetic porewaters for each study are provided in Table 27. The model approach defined in this work was also applied to the published data. The model predictions, shown in Figure 52, can explain relatively well the sorption of U(VI) onto all clayey rocks. In Figure 52b, the model prediction slightly overestimates the experimental data reported by Marques Fernandes et al. (2015), although, for the other studies, the model predictions are within the uncertainty of the modelling parameters (Figure 52a,c). Based on this model exercise, we can validate that considering only one carbonate surface complexation (≡S s OUO 2 (CO 3 ) 2 3-) for illite, together with the surface reactions for both illite and montmorillonite (Table 19) is enough to quantify the behaviour of such complex systems as natural clay rocks.

Summary

The aim of the present work was to propose a model approach able to describe the retention of U(VI) on Callovo-Oxfordian formation as a basis for the study in temperature. The 2SPNE SC/CE model proposed by Bradbury and Baeyens (1997) was used considering the simplification already proposed in the literature (Bradbury and Baeyens, 1997, 2005c, 2011;Chen et al., 2014;Marques Fernandes et al., 2015): Callovo-Oxfordian clay fraction governs the retention of U(VI), where we consider only illite and interstratified I/S layers, I/S is a mixture of illite and smectite and montmorillonite are used as a proxy of smectite.

The sorption experimental data were performed for illite in order to better understand the impact of the formation of carbonato complexes on the retention of U(VI), taking as basis the model for montmorillonite already largely studied (Bradbury and Baeyens, 1999;Marques Fernandes et al., 2012;[START_REF] Tournassat | Cation Exchange Selectivity Coefficient Values on Smectite and Mixed-Layer Illite/Smectite Minerals[END_REF]Tournassat et al., , 2018)). The results show that the model can explain relatively well the experimental data obtained with illite in a range of experimental conditions (pH = 6-7.3: pCO2 = 10 -3.5 and 10 -2.1 ; in 0.1 M of NaNO3 and synthetic COx porewater) indicating that a carbonato complex is sorbed. The constant for the main complex (≡S s OUO 2 (CO 3 ) 2 3-) was fitted and shown to be higher than the one proposed for montmorillonite. The model of sorption isotherms of U(VI) also indicates that this surface complex is the main one responsible for the retention of U(VI), indicating that the impact of ≡S s,w OUO 2 (CO 3 ) -can be considered negligible.

The proposed model explains relatively well the behaviour of U(VI) in COx clay fraction, although some divergence was observed in the absence of carbonates. Considering that the model works in Callovo-Oxfordian system, no changes of the parameters were done, and we consider at this stage that the following assumption is valid: I/S = 50% illite + 50% montmorillonite.

Finally, the proposed model was successfully applied to natural samples of Callovo-Oxfordian formation (and other clay rocks, such Opalinus clay by Amayri et al., 2016;Joseph et al., 2013;Marques Fernandes et al., 2015 andBoda clay by Marques Fernandes et al., 2015). Our results indicate that the clay fraction effectively governs the retention of U(VI). It is important to note that the model considers an important number of parameters, such as cation exchange reactions, protolysis reactions and surface complexation reactions for illite and montmorillonite. Furthermore, it is difficult to assess the uncertainty of the parameters on the model prediction. In our approach, we showed that the uncertainty of the surface carbonato complex is relatively high.

To predict the sorption behaviour of U(VI) in Callovo-Oxfordian formation system at 20ºC we propose the surface parameters provided in Table 28.

We will use this proposed model approach to predict the behaviour of U(VI) in Callovo-Oxfordian formation at 20ºC as a basis for the following chapter, where we will evaluate the impact of temperature on the retention of U(VI).

Introduction

At this stage of the work, we already know that the temperature influences the speciation of U(VI) in calcium rich solutions, specifically that temperature does not favour the formation of CaUO2(CO3)3 2- whereas it has no effect on Ca2UO2(CO3)3(aq) species formation. In this chapter we ask if the increase of temperature can also affect the retention of U(VI) on Callovo-Oxfordian formation.

One of the answers for this question is given in Figure 53 (see symbols). The experimental data shown in this figure corresponds to the experiments performed with the COx claystone natural sample (the experimental results are detailed in Section 4.2.4). The results indicate an increase of the sorption of U(VI) with a temperature increase. Such effect was already observed in Opalinus clay conditions for U(VI) (Joseph, 2013) and Np(V) [START_REF] Fröhlich | Influence of Humic Acid on neptunium(V) Sorption and Diffusion in Opalinus Clay[END_REF], or for U(VI) in granite systems [START_REF] Jin | Surface Complexation Modeling of U(VI) Adsorption on Granite at Ambient/elevated Temperature: Experimental and XPS Study[END_REF]. In conclusion, the sorption of U(VI), independently of the material, is commonly favoured by the increase of the temperature.

The increase of U(VI) sorption can be predicted by modelling the retention of U(VI)/COx system using the model proposed in Chapter 3 at 20 ºC, including the thermodynamic data for the Ca-U(VI)-CO3 ternary complexes, given in Chapter 2. The presence of Ca-U(VI)-CO3 aqueous complexes may explain the weak retention of U(VI) at 20°C while with an increase in temperature these aqueous species are less stable in solution favouring the sorption of U(VI). However, the model does not fully predict the U(VI) sorption increase in the COx system (red solid-line in Figure 53). We must therefore consider other parameters (such as enthalpy) that will allow us to fully understand the retention mechanisms of U(VI) with the increase of temperature.

In the last years, it has been shown that the mineral assemblage of Callovo-Oxfordian formation does not change with an increase of temperature of up to 80 ºC [START_REF]Dossier 2005 Argile -Synthèse Evaluation de La Faisabilité Du Stockage Géologique En Formation Argileuse[END_REF] and only the chemical composition of the porewater changes accordingly (see Chapter 1). Therefore, the question is how the surface reactions constants are affected with a temperature increase, in the case for instance of surface protolysis reactions, ionic exchange reactions and surface complexation reactions. This question will be discussed in the following section, taking into account the published studies.

Thermodynamic parameters for sorption processes

The thermodynamic parameters (enthalpy and entropy) of sorption reactions have been studied for a variety of metal ions in different systems such as oxides, model phases such as silica and alumina, and even in clay minerals (kaolinite, smectite, illite) [START_REF] Fokkink | Temperature Dependance of the Electrical Double Layer on Oxides: Rutile and Hematite[END_REF][START_REF] Fokkink | Temperature Dependence of Cadmium Adsorption on Oxides. I. Experimental Observations and Model Analysis[END_REF][START_REF] Brady | Silica Surface Chemistry at Elevated Temperatures[END_REF][START_REF] Brady | Alumina Surface Chemistry at 25, 40, and 60 °C[END_REF][START_REF] Brady | Surface Chemistry and Silicate Dissolution at Elevated Temperatures[END_REF][START_REF] Angove | The Influence of Temperature on the Adsorption of Cadmium ( II ) and Cobalt ( II ) on Kaolinite[END_REF]Halter, 1999;[START_REF] Shahwan | Temperature Effects in Barium Sorption on Natural Kaolinite and Chlorite-Illite Clays[END_REF][START_REF] Tertre | Experimental Sorption of Ni 2+ , Cs + and Ln 3+ onto a Montmorillonite up to 150°C[END_REF]Tertre et al., , 2006;;[START_REF] Bauer | Influence of Temperature on Sorption of Europium onto Smectite: The Role of Organic Contaminants[END_REF][START_REF] Duc | Temperature Effect on the Flávia Maia | IMT Atlantique Nantes Acid-Base Behaviour of Na-Montmorillonite[END_REF][START_REF] Yang | Temperature Effect on U(VI) Sorption onto Na-Bentonite[END_REF]Gao et al., 2015;[START_REF] Jin | Surface Complexation Modeling of U(VI) Adsorption on Granite at Ambient/elevated Temperature: Experimental and XPS Study[END_REF].

Protolysis reactions:

In some oxides, such as silica and alumina, the behaviour of the surface charge and the respective sorption of metal cations can change with temperature. [START_REF] Brady | Silica Surface Chemistry at Elevated Temperatures[END_REF][START_REF] Brady | Alumina Surface Chemistry at 25, 40, and 60 °C[END_REF]) studied the temperature effect on SiO2 and α-Al2O3 reactivity between 25 and 60ºC. Both studies intended to determine the enthalpy of reactions of the most relevant sorption sites in such minerals. The authors concluded that the protonation-deprotonation surface sites in both silica and alumina could be affected by the increase of temperature, which was also observed by Halter (1999). These sites are pH-dependent which means that the pH-dependency increases with the increase of temperature. This observation was corroborated by studies on the dissolution of aluminosilicates by [START_REF] Brady | Surface Chemistry and Silicate Dissolution at Elevated Temperatures[END_REF]. Duc and co-authors (2008) studied the effect of temperature on surface protolysis reactions of montmorillonite, which was considered a combination of alumina and silica layers, in more complex systems such as the clay minerals. In contrast to what was observed by [START_REF] Brady | Silica Surface Chemistry at Elevated Temperatures[END_REF][START_REF] Brady | Alumina Surface Chemistry at 25, 40, and 60 °C[END_REF] and [START_REF] Brady | Surface Chemistry and Silicate Dissolution at Elevated Temperatures[END_REF], the protonation-deprotonation at the edge surface sites is not temperature-dependent on montmorillonite; in particular, the PZNPC (point of zero net proton charge) did not change significantly with the increase of temperature. This effect was also observed by Tertre et al. (2006a).

Ionic exchange and surface complexation reactions:

The temperature effect on the adsorption of metal cations has also been studied for different clay minerals in a range of 20 to 150 ºC (e.g. Brady et al., 1992Angove et al., 1997, 1998;[START_REF] Bauer | Influence of Temperature on Sorption of Europium onto Smectite: The Role of Organic Contaminants[END_REF][START_REF] Shahwan | Temperature Effects in Barium Sorption on Natural Kaolinite and Chlorite-Illite Clays[END_REF][START_REF] Tertre | Experimental Sorption of Ni 2+ , Cs + and Ln 3+ onto a Montmorillonite up to 150°C[END_REF]Tertre et al., , 2006;;[START_REF] Chang | Sorption of Ni(II) on Na-Rectorite from Aqueous Solution: Effect of pH, Ionic Strength and Temperature[END_REF][START_REF] Donat | The Removal of Uranium (VI) from Aqueous Solutions onto Natural Sepiolite[END_REF][START_REF] Sun | The Retention of Uranium and Europium onto Sepiolite Investigated by Macroscopic, Spectroscopic and Modeling Techniques[END_REF][START_REF] Yang | Temperature Effect on U(VI) Sorption onto Na-Bentonite[END_REF]. In general, all studies indicate temperature-dependency of the adsorption of metal cations in clay mineral systems. The adsorption of the metal cations increases with temperature, indicating an endothermic process. This however depends on the mechanism of retention considered, for example [START_REF] Tertre | Experimental Sorption of Ni 2+ , Cs + and Ln 3+ onto a Montmorillonite up to 150°C[END_REF]Tertre et al. ( , 2006) ) showed that the temperature effect depends on the surface complexation mechanisms, which slightly increase when the temperature increases, whereas the influence of the cation exchange is negligible.

Even if a general trend, indicating the endothermic character of the metal cations sorption, can be perceived from most studies it cannot be generalised. For example, [START_REF] Bauer | Influence of Temperature on Sorption of Europium onto Smectite: The Role of Organic Contaminants[END_REF] show that the effect of temperature on the adsorption of Eu(III) on smectite was negligible.

Conceptual approach

The literature review allowed us to conclude that a temperature effect and the surface reactivity cannot be excluded. However, to the best of our knowledge, there are no enthalpy values which can be directly applied to the surface parameters describing U(VI)/Callovo-Oxfordian formation at 20 ºC (see Table 28 in Chapter 3).

Simplifications of the model in temperature must be done in order to decrease the number of parameters to study. Based on the literature review we can carefully assume that:

-Since the protolysis reactions are not temperature-dependent for montmorillonite (Tertre et al., 2006a;[START_REF] Duc | Temperature Effect on the Flávia Maia | IMT Atlantique Nantes Acid-Base Behaviour of Na-Montmorillonite[END_REF]) we can apply the same assumption for illite; -The temperature effect is also negligible for the cationic exchange reactions [START_REF] Tertre | Experimental Sorption of Ni 2+ , Cs + and Ln 3+ onto a Montmorillonite up to 150°C[END_REF](Tertre et al., , 2006)); Despite these simplifications, there are still an important number of surface complexation constants to consider in the model and this remains a challenge, in which a large number of experimental data should be acquired in order to allow all the species to be formed with different ranges of existence.

The complexity of such experimental and modelling approach was also raised in the work of Jin and co-workers (2016) while evaluating the main mechanisms involved in the sorption of U(VI) on granites under different conditions (pH, ionic strength, initial concentration of U(VI) and temperature). The authors also observed an increase in U(VI) sorption in temperature, proposing enthalpy values for the different surface complexation reactions involved. The determination of the thermodynamic data was based on the model fit constraining the system to three surface complexation reactions at constant pH (4.5, 5.2 and 6.3). Considering the complexity of the granite composition, the generalised model approach described the U(VI) adsorption at high concentrations of U(VI) (> 10 -5 M) in equilibrium. These authors also aimed to understand the influence of carbonates on U(VI)/granite systems, but their model was not able to interpret the experimental data.

The determination of enthalpy values for surface species could be helped by analysing the thermodynamic data known for the aqueous system. A good correlation between surface complexation constants and the corresponding species in aqueous phase at 25ºC was already shown for the parameters of the 2SPNE SC/CE model (Bradbury and Baeyens, 2009). We could therefore expect such good correlation in temperature. [START_REF] Brady | Silica Surface Chemistry at Elevated Temperatures[END_REF][START_REF] Brady | Alumina Surface Chemistry at 25, 40, and 60 °C[END_REF] showed that the adsorption of Cd(II) and Pb(II) is favoured with an increase of temperature for silica or alumina and there was a good correlation between the aqueous complexation and the respective surface complexation reactions. Their results indicated that the enthalpy of hydrolysis of metal cations is positively correlated with the enthalpy of silica surface, but weakly correlated with the enthalpy of alumina surface.

Thus, the model approach proposed in this study will take into account the correlation between the aqueous complexation and surface complexation thermodynamic data, in addition to the two other assumptions mentioned above (temperature effect negligible for the protolysis reactions and cation exchange reactions).

Effect of the temperature on the adsorption processes of uranium on clays and clay minerals

The temperature changes of the system affect not only the chemical composition and properties of the solution but also the surface reactivity of the solids. The overall system changes will have implications on the uranium's behaviour on Callovo-Oxfordian formation. We evaluated the effect of the temperature on its behaviour through the analysis of sorption experiments carried out in different mineral assemblages at different conditions and different temperatures, together with modelling exercises.

Experimental approach

As in Chapter 3, we started first with the analysis of a single mineral, illite, followed by an increase of the complexity of the system, through the analysis of the clay fraction (illite and interstratified I/S layers) and the natural samples (clay minerals, carbonates and silicates) of the Callovo-Oxfordian formation, measuring and assessing the effect of pH and carbonate contents of the solution at different temperatures. The experimental cases are described in Table 29.

The experimental set-up for both the sorption edge and isotherm experiments follows the same procedure as described in Chapter 2 (see Figure 6) using the chemical conditions as described in Table 30. It should be noted that all solutions were prepared at room temperature (approx. 20ºC) and heated during the period of equilibration using the incubator adjusted to the temperature of interest. The Flávia Maia | IMT Atlantique Nantes material and methods used for the sorption experiments are described in Appendix A, and the details of the experimental results are provided in Appendix C.5.

S/L (g/L) 2.0 0.5-1.0 2 2 2 2 2-10 [U(VI)]initial (M)
10 -7 10 -7 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5 10 -8 -2•10 -5 Equilibration (days) 3 3 3 3 3 3 3

Model Approach

The model approach that we propose for studying the temperature effect on the retention of U(VI) considers that the ionic exchange reactions [START_REF] Tertre | Experimental Sorption of Ni 2+ , Cs + and Ln 3+ onto a Montmorillonite up to 150°C[END_REF](Tertre et al., , 2006) ) and the protonation/deprotonation reactions (Tertre et al, 2006a;[START_REF] Duc | Temperature Effect on the Flávia Maia | IMT Atlantique Nantes Acid-Base Behaviour of Na-Montmorillonite[END_REF] are not influenced by the temperature. The thermodynamic data describing the surface complexation reaction with U(VI) are unknown and will be derived from the experimental data. Note that the two surface carbonato for illite, which do not play a major role at 20°C for the studied experimental conditions (see Section 3.4.2), will not be considered in this study. The model will be developed firstly using the U(VI)/illite system in order to propose a modelling strategy which will be further used, in a predictive way, to the COx clay fraction and claystone natural sample.

We will use the model approach 2SPNE SC/CE proposed in Chapter 3 at 20 ºC (Model II) using the thermodynamic data of the aqueous species (Appendix E), i.e. no enthalpy values are considered for the surface species. This way we will be able to evaluate from the beginning if the surface species on illite (and after on montmorillonite) are, or not, controlled by enthalpy.

As mentioned previously, there is a correlation between complexation in solution and complexation at the surface (Figure 54). Considering the important number of surface reactions to be considered, we decided to assume in a second approach a correlation factor of 1 for the known enthalpies, from the aqueous species complexation reactions to the corresponding surface complexation reaction. Depending on the results, the enthalpy values will be fitted in order to obtain a good agreement between the experiment and the model, aiming to obtain a correlation between complexation and surface complexation constants.

i Uncertainty not available We will describe next the conceptual approach used to link the hydrolysis and carbonato surface complexation, using the corresponding aqueous species, as described in Bradbury and Baeyens (2009).

The hydrolysis constants ( OH Kx values) for the metal Me z are given in the literature for aqueous hydrolysis reactions. The general form can be written as follows:

Me z + xH 2 O ⇌Me(OH) x z-x +xH +

Reaction 21

In Reaction 21, x is an integer > 1, and Me(OH) x (z-x) is a general form corresponding to the hydrolysed species.

The reaction between aqueous hydrolysed species and amphoteric surface hydroxyl sorption sites can be written in a generalised form as:

≡S s OH+Me z + yH 2 O ⇌ ≡S s OMe(OH) y z-(y+1) +(y+1)H + Reaction 22
In Reaction 22, y is an integer, and, when y = 0 the surface complex is ≡S s OMe z-1 . The corresponding surface stability constant is expressed as S Ky, through the standard mass action relationship, and the general form of the surface species is ≡S s OMe(OH) y z-(y+1) . In the approach proposed by Bradbury and Baeyens ( 2009), OH Kx is correlated to S Ky-1.

In the case of the carbonation constants ( CO3 Kx values) for Me z the action mass law can be expressed in a general form as:

Me z + xCO 3 2-⇌Me(CO 3 ) x (z-2x)

Reaction 23

In Reaction 23, x is also an integer > 1. Taking the Reaction 22 and assuming y = 0, and that the reaction between aqueous carbonate species and amphoteric surface hydroxyl sorption sites exists, the mass action reaction of the carbonation of surface can follow the same principle described above and it can then be written as:

≡S s OH+Me z + yCO 3 2-⇌ ≡S s OMe(CO 3 ) y (z-2y) + H + Reaction 24
The corresponding surface stability constant is expressed as S Kx, via a standard mass action relationship, and the surface species is ≡S s OMe(CO 3 ) y (z-2y) . If we follow the approach of Bradbury and Baeyens (2009), the surface species ≡S s OMe(CO 3 ) y (z-2y) must correspond to the aqueous mixed complex HOMe(CO 3 ) y (z-2y-1) . However, such species are not proposed in the thermodynamic database, and as such the surface enthalpy values will be directly evaluated by fitting, if necessary.

In our study, Me z is UO2 2+ and the main hydrolysis and carbonate aqueous and surface complexes and their respective complexation reactions are described in Table 31. The ionic exchange reactions and protonation/deprotonation reactions are also included in the model while enthalpy values are not included (see Table 28).

U(VI)/illite system in temperature up to 80 ºC

We performed sorption edges and isotherms experiments to study the effect of temperature on U(VI) sorption onto illite at 60 and 80 ºC, in the absence or in the presence of carbonates in equilibrium with different pCO2(g) atmospheres. The experimental data will be considered in order to propose a modelling strategy which will be further used to predict the U(VI) sorption on the COx clay fraction (in order to assess whether it can be used also with smectite, i.e. montmorillonite) and COx natural samples (to evaluate its applicability to a natural complex mixture with additional mineral phases).

In the absence of carbonate

Experimental data:

In Figure 55, we show the distribution coefficient values of U(VI) obtained as a function of the pH at 80 ºC in the absence of carbonates (in inert atmosphere). In comparison with the results obtained at 20 ºC, the experimental behaviour of U(VI) does not change at neutral conditions, however it is slightly higher at low pH. At pH below 5, the KD increases about 1 order of magnitude with the increase of temperature to 80 ºC.

The sorption isotherms of U(VI) at 60 and 80 ºC are compared with the sorption isotherms performed at 20 ºC (Figure 56). At the end of the experiments the pH measured reached a value of 6.4 (±0.5). Note that the alkalinity titration measurements indicate that in these experiments there is a possible contamination of carbonates in the glove-box. This issue was already observed when we performed the experiments at 20 ºC. The corresponding pCO2 values were calculated and are indicated in Table 32. The uncertainties of the alkalinity measurements are relatively high considering the low content of carbonates measured. The contamination was taken into account in the model without considering at this point an enthalpy value for the carbonate surface species. Table 32. pH and alkalinity measured in equilibrium for the sorption isotherms on illite in inert atmosphere at different temperatures.

Temperature pHfinal Alkalinity (meq/L) pCO2 calculated i 20 ºC 6.4 (± 0.2) 0.08 (± 0.03) -2.7/-3.0 60 ºC 6.4 (± 0.5) 0.08 (± 0.03) -2.5/-2.8 80 ºC 6.5 (± 0.1) 0.08 (± 0.03) -2.5/-2.9

As it was observed for the sorption edge experiments, the experimental results show no significant change in sorption when the temperature increases; a slight increase is nonetheless observed at concentrations below 10 -8.5 M for temperatures up to 80 ºC.

Modelling results:

Firstly, we performed blind predictions to assess the influence of temperature on aqueous speciation (Figure 57). These results show that the increase of temperature "shifts" the aqueous species distribution to a lower pH. For example, at neutral pH the predominant aqueous species at 80 ºC is UO2(OH)2 instead of UO2OH + at 20 ºC. We reason that this must affect the surface complexation reaction.

As mentioned previously, in a first approach, we did not consider any enthalpy values in the geochemical model for the surface species (Model II). In a second approach, the enthalpy values for the aqueous species were directly applied to the corresponding surface complexation species (i.e. considering a correlation factor of 1) (Model IV, Table 33). The results are shown in Figure 58. The model approach that does not consider a temperature effect on the surface reaction (Model II) cannot explain the experimental data at 80 ºC: the model predicts a strong decrease of the retention as a function of temperature, whereas a slight effect is observed experimentally at lower pH. This can be explained by the strong influence of the temperature on the formation of the hydrolysis species (as it was shown in Figure 57), which increases the competition between the formation of aqueous species and the surface complexation.

When the enthalpy values for the hydrolysis surface complexation reactions are considered (Model IV), the model approach can effectively reproduce an increase of the sorption of U(VI) as a function of pH at 80 ºC (Figure 58b), indicating that the hydrolysis surface complexation is also favoured in temperature. However, this approach highly overestimates the sorption of U(VI) at pH below 7.5 at 80 ºC. Furthermore, the model can no longer reproduce the experimental data at 20 ºC for pH values between 6 and 8 (Figure 58a). The same effect is observed for the sorption isotherms of U(VI) at different temperatures (20, 60 and 80 ºC) (Figure 59a, b and c, respectively). 33). Error bars that are not visible are smaller than the symbol size.

From these results, we conclude that the slight increase of U(VI) sorption with the increase of temperature can be quantitatively explained by considering a positive effect of temperature on the surface complexation reactions, which can "compensate" the same result predicted in solution thus resulting in a "null" effect. This positive effect cannot be explained by applying correlation factor of 1 for the enthalpy values (following the correlation of the stability constants for the surface and aqueous species proposed by Bradbury and Baeyens (2009)) and they must be fitted. 33). Error bars that are not visible are smaller than the symbol size.

This fitting approach was done manually since PEST is not the most suitable tool, as it does not provide reliable fitting parameters as a consequence of the high number of equilibria reactions and the limited number of experimental points. The procedure to determine the enthalpy values for the surface complexation reactions relies on five criteria:

I. The surface speciation shown in Figure 60 was used as a guide for finding the best fit for: 1) ≡S s OUO 2 + in pH < 5; 2) ≡S s OUO 2 OH between pH 5 and 6; and 3) ≡S s OUO 2 O(OH) 2 -at pH > 6 together with the ≡S s OUO 2 (OH) 3 2-. The iterative procedure was performed considering both the sorption edge and isotherm data. Note that since we only have the sorption isotherm of U(VI) at pH of 6.4-6.5, the constants of ≡S s OUO 2 (OH) 2 -and ≡S s OUO 2 (OH) 3 2-will be relatively more reliable; II. The surface constants obtained in temperature must be correlated with those describing the formation of the corresponding aqueous species; III. The same enthalpy values are fixed for strong and weak sites when the surface species are similar (i.e. ≡S s,w1 OUO 2 + and ≡S s,w1 OUO 2 OH);

IV. According to the van't Hoff equation, the complexation constants at different temperatures can be expressed as Equation 3i ;

Since the surface standard constants for illite are known (log10 OH Kxº(25 ºC), see Table 31), the Equation 3 can be readjusted to determine the enthalpy values for each surface species (Equation 10).

log 10 K s °(T) -log 10 K s °(T 0 ) = - ∆ r H°(T 0 ) Rln10 ( 1 T - 1 T 0 )
Equation 10

V. Applying the enthalpy values should not significantly modify the modelling at 20°C with respect to that presented in Chapter 3. 33.

The results of the fitting process considering the five previous criteria are presented in Table 34 and shown in Figure 61. The agreement between the model and the experimental data is good for both the sorption edge and isotherms data. For the sorption edges at 20 and 80ºC and the sorption isotherms at 20ºC the uncertainty of the surface constants was estimated to be ± 0.2 log units, whereas for the i As reminder of Equation 2:

log 10 K s °(T)=log 10 K s °(T 0 )- ∆ r H°(T 0 ) Rln10 ( 1 T - 1 T 0 )
Impact Of The Temperature On The Sorption of U(VI) in The Callovo-Oxfordian Formation sorption isotherms at 60 and 80 ºC the uncertainty was found to be ± 0.5 log units (in order to cover maximum number of experimental data). The correlation between the surface and the aqueous complexation constants obtained at different temperatures in the log-log scale is shown in Figure 62. This correlation is similar to the one shown in Figure 54 (compare solid-line for both Figure 54 and Figure 62), indicating that the second criterion is respected. The dependency of constants with the temperature using the formalism described in Equation 10 is given in Figure 63. Despite the fact that the uncertainties for the constants are significantly high, especially for the surface species ≡S s OUO 2 + and ≡S s OUO 2 OH, globally the expected linear tendency is observed. The enthalpy values determined by the fitting are shown in Table 35. The final results are given in Figure 64 using the enthalpy values for the hydrolysis surface complexes given Table 35. The agreement between the calculated and measured concentration of U(VI) is good if we consider the sorption isotherm data. For the sorption edge data, the model approach (Model V) is not able to fully reproduce the behaviour of U(VI) as a function of pH, although an improvement is observed comparatively to when using Model IV, which considers the enthalpy values with a correlation factor of 1 (see Figure 58). In summary, the temperature effect on the retention of U(VI) on illite in the absence of carbonates can be explained by considering an endothermic effect to describe the surface complexation reactions between U(VI) and U(VI) hydrolysis species with the strong and weak sites present at the illite surface. The next step is to apply Model V to conditions where carbonato complexes are formed to study their behaviour in temperature. These data will also allow us to assess the potential impact of the contamination of carbonates, found in the experiments performed under inert atmosphere, on the parameters proposed in Table 35.

In the presence of carbonate Experimental data:

We performed the sorption edge experiments in atmospheric pCO2 at 80 ºC in 0.1 M of NaCl and the initial concentration of U(VI) was set at 3.3•10 -8 M (Figure 65). The alkalinity titration and pH measurements were done before and after the equilibrium and the results are provided in Table 36. The results indicate that the alkalinity does not change significantly with the increase of temperature, whereas the pH slightly decreases with the increase of temperature. The corresponding calculated pCO2 shows that the equilibrium is not completely reached for most pH values. The sorption of U(VI) appears similar at 80 °C for the experiments performed in the presence or absence of carbonates when the pH is below 7.5 (Figure 65). For pH > 7.5, a decrease of the sorption is observed resulting from the formation of strong complexes that stabilise U(VI) in solution. This indicates a weak influence from surface speciation on the retention at 80°C. In comparison, at 20 °C and above pH 6 we observed a significant difference on the retention of U(VI) between the experiments performed in inert atmosphere and those performed in atmospheric pCO2 (see Figure 37 in Chapter 3). The sorption isotherms of U(VI) in equilibrium with different pCO2 at different temperatures are shown in Figure 66. The alkalinity of the solution is the same at pCO2 of 1 or 5% of CO2(g) but the pH of equilibrium is lower at 5% of CO2(g). The calculations indicate a water/gas system close to equilibrium (taking into account the uncertainty of the alkalinity titration measurements, see Table 37). A qualitative analysis indicates that temperature changes do not have a strong effect, and when observed this effect is slightly positive (i.e. there is a increase on the retention of U(VI) with the increase of temperature). 2.4(±0.9) -1.5/-1.9

Modelling Results:

The behaviour of U(VI) in solution was first studied by blind modelling prediction for atmospheric pCO2. The results are presented in Figure 67 at 20 ºC (solid-lines) and 80°C (dashed-lines). An important observation is that the presence of uranyl-carbonate aqueous complexes is much less relevant for pH < 7 at 80 ºC as compared to 20 °C. The temperature increasing favours the formation of hydrolysis species in detriment of carbonate ones. For example, at pH 7 the proportion of carbonate complexes changes from 60% to 10% for 20 and 80 °C, respectively. Above pH 7 the U(VI) speciation is governed by the UO2(CO3)3 4-at 20 ºC and UO2(OH)3 -at 80 ºC.

The distribution of the U(VI) speciation can explain qualitatively the sorption experiments, indicating that at pH >7 the retention decreases (as at 20°C) due to the formation of U(VI)-CO3 complexes, even if that retention is small (compare orange and blue dashed-lines in Figure 67 with symbols in Figure 68). Below pH 7, the presence of carbonates does not influence the retention since U(VI) speciation is dominated by the hydrolysis species (see green dashed-line in Figure 67 and symbols in Figure 68). Based on these results, we propose to model firstly the data considering only a temperature effect on the surface hydrolysis complexes, i.e. Model V (see above Table 35). These modelling results correspond to the red solid-lines in Figure 68 and Figure 69. Overall, the agreement between the calculation (modelling) and the experimental data is acceptable. The presence of trace amounts of carbonates for the experiments performed in inert atmosphere does not affect the modelling results. The model can reproduce well the pH-edge for atmospheric pCO2 but it overestimates the retention for the isotherms obtained in temperature in the presence of CO2, particularly at 80 °C. The analysis of the surface speciation below pH 7 (thus including the isotherm in concentration) indicates that the retention is governed by the surface hydrolysis species. At this point and in agreement with the qualitative analysis of the data presented in Figure 66, we can conclude that the surface carbonate complexes do not play any role on the sorption of U(VI) at 60 and 80°C and there is consequently no reason to consider a temperature effect on the formation of such surface complexes.

In conclusion, we propose from here to use Model V and to assess its ability to explain sorption data in temperature in the context of the Callovo-Oxfordian clay fraction and claystone natural samples.

U(VI)/ Callovo-Oxfordian clay fraction system at 80ºC

In the present work we have also studied the increase in complexity of the mineral assemblage in temperature. As mentioned in the previous chapter, at 20 ºC the sorption of U(VI) slightly changes in the presence of the different clay minerals. In this section, we study the behaviour of U(VI) on clay fraction of Callovo-Oxfordian formation at 80ºC for different solution compositions (see Table 30), focusing on the measure of sorption isotherms in concentration of U(VI) only.

The experimental results on the sorption isotherms of U(VI) at 80 ºC in the absence or presence of carbonates (in equilibrium with 5% of CO2(g)) are shown in Figure 70a andb. The pH and alkalinity measurements performed at the end of the experiments are provided in Table 38. The results indicate that there is a contamination of carbonates in the "inert" system. As compared to the previous experiments (U(VI)/illite system, Table 32) where the alkalinity was relatively small with high uncertainties, the alkalinity here appears relevant, and as such it was considered in the modelling analysis. As already observed at 20 ºC, in the absence of carbonates the sorption of U(VI) decreases for the clay fraction (Figure 70a) as compared to illite, whereas in the presence of carbonates, no significant differences are observed between the illite system or the clay fraction system (Figure 70b). Ar(g) 7.2 (± 0.1) 0.22 (± 0.1) -2.8/-3.2 5 % of CO2(g) 7.2 (± 0.02) 2.2 (± 0.9) -1.8/-2.2

By comparing the experimental data at 20 and 80ºC in the presence or absence of carbonates we can conclude that the temperature effect is not significant, as it was previously observed for illite (Figure 71a,b). As reported by Bradbury and Baeyens (2009), illite and montmorillonite are reactively similar. Thus, the same thermodynamic considerations were considered for both clay minerals. The thermodynamic parameters that we considered to describe the retention of U(VI) on the Callovo-Oxfordian clay fraction are given in Table 39. Overall, the model can explain relatively well the experimental results, especially in the presence of carbonates at 20 and 80ºC. In the absence of carbonates, the model approach overestimates the sorption for tracer concentrations of U(VI) (< 10 -7 M).

In conclusion, the use of the enthalpy values for illite appears applicable to montmorillonite. In the conditions of interest for the Callovo-Oxfordian formation regarding pH and alkalinity, the temperature increase will favour the predominance of hydrolysis species with respect to carbonato complexes, both in solution and at the surface. In the following section we will assess and discuss the impact of the presence of ternary complexes on retention in temperature.

Impact of Ca-U(VI)-CO 3 on the retention of U(VI) at 80 ºC

We performed the sorption isotherms experiments in concentration on illite, COx clay fraction and natural samples at 80 ºC in equilibrium with synthetic COx porewater. The sorption experiments were performed in closed atmosphere with 1% of CO2(g) for experiments at 20 ºC and 5% of CO2(g) for experiments at 80 ºC and the synthetic porewater was prepared following the composition given in Table A. 2. The pH of equilibrium and alkalinity measurements for each case are provided in the Table 40. Clay fraction 20 ºC 1% CO2(g) 7.1 (± 0.02) 2.4 (± 0.9) -2.3/-1.9 80 ºC 5% CO2(g) 7.1 (± 0.04) 3.4 (± 1.3) -1.8/-1.5

Claystone 20 ºC 1% CO2(g) 7.5 (± 0.01) 3.3 (± 1.3) -2.7/-2.2 80 ºC 5% CO2(g) 7.0 (± 0.01) 1.4 (± 0.8) -2.3/-1.7

The experimental results show that the temperature increase leads to an increase in the sorption capacity of U(VI) in all cases. The KD of illite increases from 37 (± 9) to 153 (± 47) mL/g; for COx clay fraction it increases from 78 (± 39) to 157 (± 50) mL/g and for COx natural samples, the KD increases from 0.5 (± 1) to 263 (± 31) mL/g (see symbols in Figure 72). The temperature effect is more pronounced for COx natural samples, where the sorption capacity increases more than two orders of magnitude. We found that the difference in pH/alkalinity between the two temperatures is the highest for the Callovo-Oxfordian system.

We applied the model approach proposed in this work to the different solid phases (illite, clay fraction and claystone natural samples), and the results are shown in Figure 72 (solid-and dashed-lines). This model approach can explain the tendency for the presence of an endothermic character of U(VI) sorption when the temperature increases, although with a slight overestimation for U(VI)/illite system and U(VI)/clay fraction at 80ºC. In conclusion, the model that we propose, which considers the enthalpy values for the surface complexation reactions for both illite and montmorillonite, can predict relatively well the behaviour of U(VI) in different systems in the presence of Ca-U(VI)-CO3 ternary complexes (Figure 73). A summary of the most relevant aqueous and surface complexation reactions is given in Table 41 for the U(VI)/Callovo-Oxfordian system. Although the formation of ternary complexes is not favoured in temperature these are still the dominant species at 80 °C in solution. In contrast, at the surface at 20 °C the sorption is dominated by the carbonate surface complex species whereas at 80°C the surface hydrolysis species are dominant. Moreover, the clay fraction, including both illite and illite from I/S interstratified clay, represents the main reactive surface at both temperatures.

It is interesting to note that Joseph and co-workers (2013) suggested that this increase in retention with temperature was due to the formation of U(VI) colloids which are formed when the temperature increases. In the present study, we did not evaluate the possible formation of colloids; although, based on ICP-MS analysis of the blank solutions we do not have evidences of their occurrence. However, we can expect a possible precipitation of U(VI) phases based on the evaluation of the geochemical calculations using the thermodynamic database available.

Summary

The behaviour of U(VI) sorption in temperature was studied from the simplest clay system (illite) to the most complex (with Callovo-Oxfordian claystone mineralogical assemblage), where the pH, the pCO2 and solution composition were changed.

As already observed for the different conditions presented in Chapter 2, the control of the pH/alkalinity at a given pCO2 is not straightforward, affecting more the experiments performed as a function of the pH. For modelling purposes, in the experiments at constant pH, the pCO2 was never fixed, it was instead freely controlled by the pH and alkalinity measurements, even though we are aware that the uncertainty of the alkalinity analysis is usually high and small changes of the alkalinity values can lead to different results. Although, we recommend that the pCO2 should be fixed at pH > 7.5.

A main conclusion of this study is that the effect of temperature on the sorption of U(VI) is almost negligible for ionic strength of 0.1 M, using NaNO3 or NaCl as electrolyte background, for both edgeand planar-surfaces of illite. This is in contradiction with the strong effect of the temperature expected in solution, based on the thermodynamic data available (ThermoChimie database, Giffaut et al., 2014). Indeed, the formation of the aqueous complexes (specially the hydrolysis complexes) are favoured in temperature and this would then lead to a decrease of the retention.

To allow a quantitative explanation of the data, it is therefore necessary to consider also an endothermic effect of surface complexes in order to counterbalance the aqueous chemistry effect. We cannot explain the data by applying directly the enthalpy values of aqueous species to the corresponding surface species as the model always overestimates the experimental data. We therefore fitted the enthalpy values from the experimental data following a stepwise approach with constraints (taking the van't Hoff equation and the linear correlation between the aqueous and surface speciation).

We can obtain a relatively good agreement between the model and the experimental data when considering the fitted enthalpy values associated to the hydrolysis species. We do not need to consider an endothermic effect for surface carbonate species, which are not favoured with an increase in temperature, provided that the pH remains below 7. We provide a summary of the final surface parameters used for the model approach in Table 42.

Thus, the model developed for illite can apply satisfactorily to COx clay fraction and COx claystone natural samples, in both the case of absence or presence of the Ca-U(VI)-CO3 ternary complexes.

It is clear that the uncertainties associated to the enthalpy values are high and, in the present work, they were estimated using "manual" fitting, as more robust data can be obtained by increasing the number of experimental data. Nevertheless, the model reproduces the Callovo-Oxfordian conditions and can be applied to predict in situ conditions.

Chapter 5 Conclusions

The aim of this study was to have a better understanding and a better quantification of the behaviour of U(VI) on Callovo-Oxfordian (COx) formation, which will host high-level radioactive waste in France. We have focused our study on the effect of an increase in temperature of up to 80 °C on U(VI) behaviour in solution and onto mineral surfaces.

The first part of the work focused on determining the thermodynamic properties of the calcium uranyl carbonate aqueous complexes, which govern the speciation of U(VI) in the clay porewater's surroundings of the repository. For that, an experimental method was developed in controlled atmospheres to determine the stability constants of the Ca-U(VI)-CO3 ternary complexes at different temperatures: 20, 40, 60 and 80 ºC.

The main experimental challenge of this work was to control the pH/carbonates equilibrium with temperature during the batch experiments. In initial experiments we guaranteed the correct performance of the experimental set-up and evaluated the equilibrium between CO2(g) and the solutions.

We determined the stability constants of Ca-U(VI)-CO3 complexes at different temperatures by fitting a geochemical model to the experimental data. The results showed that the two aqueous complexes studied have different behaviours in temperature. The formation of the anionic complex (CaUO2(CO3)3 2-) is not favoured with an increase in temperature, while the neutral complex (Ca2UO2(CO3)3(aq)) forms independently of temperature changes.

We obtained, from these data, the thermodynamic data for the CaUO2(CO3)3 2-and Ca2UO2(CO3)3(aq). The log10 βº113 and log10 βº213 are consistent with those previously published, which provided confidence on the methodology used and the conceptual model proposed. Thus, we determined the ΔrHº of Ca-U(VI)-CO3 ternary complexes, proposing new enthalpy data for CaUO2(CO3)3 2-(27.4 ± 8.2 kJ/mol) and Ca2UO2(CO3)3(aq) (0 ± 1.8 kJ/mol). In comparison with Endrizzi and Rao (2014) enthalpy data, the anionic complex is also exothermic, although the ΔrHº is two times lower; on other hand, the neutral complex is not affected by the temperature which contradicts the value published (-47 kJ/mol). Despite the high uncertainty of ΔrHº of the Ca2UO2(CO3)3(aq), we consider that its ΔrHº is zero.

In the second part of this work we proposed a model approach to describe the retention of U(VI) on Callovo-Oxfordian formation at 20 ºC, that was further used for the study in temperature. We used the well-know 2SPNE SC/CE model to quantify the retention mechanisms of the U(VI) firstly on illite and then on Callovo-Oxfordian clay fraction system (illite and montmorillonite). Here, montmorillonite was used as proxy for smectite from the interstratified I/S minerals present in the Callovo-Oxfordian clay fraction.

Due to a lack of available data for surface-carbonato complexation reactions for illite, we performed sorption experiments at different concentrations of carbonates. As a first approach, we used the model proposed in the literature for montmorillonite and we found that it could explain the experimental data, indicating that together with U(VI)-OH surface complexes, the U(VI)-CO3 complex was also sorbed. The results also indicated that the (≡S s OUO 2 (CO 3 ) 2 3-) surface species is responsible for the sorption of U(VI) in the presence of carbonates at pH > 7. The constant of this surface species was manually fitted and the constant value was determined as 17.5 (±0.5).

Considering the new surface data for illite and the surface data published for montmorillonite, the model considering both minerals also explained the behaviour of U(VI) in Callovo-Oxfordian clay fraction systems (illite and I/S = 50% illite and 50% montmorillonite). No other parameters were needed to change in order to explain the sorption behaviour.

We have also successfully applied our model to natural samples of Callovo-Oxfordian formation and other clay rocks, such as Opalinus and Boda clay. Our results indicated that the retention of U(VI) is essentially governed by the clay fraction. Nevertheless, it was difficult to assess the uncertainty of the different parameters on the model prediction.

In the last part of the work, we discussed the temperature effect on the U(VI)/ Callovo-Oxfordian system. The retention of U(VI) in COx conditions is favoured with the increase of the temperature as already observed in the literature. However, our modelling predictions indicated that thermodynamic parameters (enthalpy data) for the aqueous speciation (including the new ΔrHº for Ca-U(VI)-CO3 aqueous complexes) and the surface data used in the model at 20 ºC were not enough to completely explain the sorption of U(VI) on the U(VI)/ Callovo-Oxfordian system at 80ºC. For this reason, we considered that enthalpy data for the surface complexation reactions was needed.

The effect of the temperature on the retention of U(VI) was negligible for illite in the absence of calcium. The formation of U(VI) hydrolysis complexes is favoured in solution with the increase of the temperature either in the absence or presence of carbonates in solution. This would lead to a general decrease of the retention due to the competition between the complexation processes occurring in solution and at the surface. To explain quantitatively the retention of U(VI), it was therefore necessary to consider an endothermic effect of surface complexes, in order to counterbalance the solution effect. Since the hydrolysis surface complexes are responsible for the retention of U(VI), their enthalpy values were determined by fitting. The new thermodynamic parameters for the hydrolysis surface complexation reactions explained the behaviour of U(VI) on illite.

On the other hand, in the presence of calcium in the system, the Ca-(VI)-CO3 aqueous species are formed, which are not favoured (or not affected) with the increase of the temperature. Considering the endothermic effect considered for the surface reactions, the retention of the U(VI) is thus favoured with the increase of the temperature.

In conclusion, the model developed for illite in temperature was applied satisfactorily to COx clay fraction and natural samples considering the presence or absence of Ca-U(VI)-CO3 ternary complexes. It was clear that the uncertainties associated to the enthalpy values are high, nevertheless, the model could reproduce the Callovo-Oxfordian conditions and can be applied to predict in situ conditions.

Perspectives

The question of the existence of the carbonato-surface complexes on either illite or montmorillonite is still unanswered. The interpretation of the sorption data for illite indicates that such complexes must be considered in the model predictions. To increase the accuracy of the proposed stability constants of carbonato-surface complexes, sorption edge experiments at 20-25 ºC should be performed in the presence of a wide range of carbonates concentration. Additionally, the interpretation of the sorption data at different temperatures revealed also a lack of knowledge concerning the enthalpy of reaction for the surface complexes for both illite and montmorillonite. As such, future experimental work, complemented by modelling, should be focused on the determination of their specific contributions for the retention of U(VI) in dependence of pH at different temperatures and different concentrations of carbonates in solution. The formed U(VI) surface complexes should be then characterised by suitable spectroscopic techniques (e.g. ATR FT-IR spectroscopy, EXAFS spectroscopy) at different temperatures.

The robustness of the model here developed for uranium's behaviour should be tested in more realistic conditions. Sorption experiments performed with Callovo-Oxfordian claystone intact samples are on progress with U(VI). Furthermore, data are being generated in order to assess the behaviour on naturally-occurring U, i.e. U distribution between the different phases and its exchangeable character. These data will derive in KD value knowing the concentration of U naturally found in the COx porewater.

Finally, although U(VI) is expected to control uranium speciation in the porewater, it is important to point out that Callovo-Oxfordian conditions are characterized by reducing conditions, where the U(IV) may also occur. U(IV) may then be an important species to consider for surface complexation reactions. To understand the mechanism of retention of U(IV) in Callovo-Oxfordian formation, we suggest that batch-type sorption experiments using U(IV) at different temperatures should be performed as a function of key parameters such as pH, pCO2, [U(IV)] and [Ca]. The nature and structure of the surface complexation for U(VI) and U(IV) should also be investigated. For example, one could use EXAFS spectroscopy to verify which are the surface species that control the sorption of U(VI) and U(IV) and clarify if there are or not uranyl carbonato surface complexes. In this way we would achieve a full understanding of the uranium's behaviour in Callovo-Oxfordian conditions and thus define a reliable KD of uranium in in situ conditions.

Appendices Appendix A: Material and Methods

A.1. ICP-MS analysis:

The ThermoScientific XSERIES 2 ICP-MS (Thermo Fisher Scientific) has been used in the present study. The ICP-MS (Inductively Coupled Plasma Mass Spectroscopy) is an analytical technique used for elemental determinations. This technique has good detection capacities, especially for the rareearth elements. It has also many advantages over other elemental analytical techniques which include: low detection limits for most of the elements; high throughput with protocol compliance; ability of handle both simple and complex matrices with low interferences due to the high-temperature of the ICP source; superior detection capability; and ability to obtain isotopic information.

A.1.1. Optimisation method

The ignition of the plasma and the optimisation of the device are controlled by the software Plas-maLab's.

After switching on the device, the number of cps of the selected elements (Li, Co, In and U) are displayed continuously, which will be useful for the optimisation of the device. The 59 Co is used to optimise the low mass ions, the 115 In for the medium mass ions, and the 238 U is used for the heavy mass ions.

The procedure of signal optimisation starts by controlling the position of the ICP torch (depth, vertical and horizontal position). Then after, the flux of gas that is expressed by L/min is set by controlling the cooling gas (> 13 L/min), the flow rate of the nebuliser, and the auxiliary gas (Ar). Once the gas flux is set, the lens system is adjusted to focus the ion beam created in the plasma. After, the amplitude of the signal is adjusted simultaneously to the bias which corresponds to tension over the entrance to the quadrupole. Here, the adjustment of the sensitivity of the signal is set. These adjustments must be done to obtain a large number of cps and a good stability of the signal of the selected elements (< 2%), and the oxide levels (< 5%), which are controlled by the Ce++/Ce and CeO/Ce ratios.

ycalc is the y-values calculated (coups), x is the x-values (concentration of the standard solutions), b is the slope of the line and c the intersection on the y-axis. The mean of all x-values is 𝑥 ̂ and the mean of all y-values is 𝑦 ̂. The position (𝑥 ̂, 𝑦 ̂) is the centroid of all points.

The slope of the line was calculated as:

𝑏 = ∑(𝑥 𝑖 -𝑥 ̂)2 ∑(𝑦 𝑖 -𝑦 ̂)2
And, the intersection as:

𝑐 = 𝑦 ̂-𝑏𝑥 ̂
The random errors of the slope and the intersection values were respectively calculated as follows:

𝑠𝑏 (𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑑𝑒𝑣𝑖𝑎𝑡𝑖𝑜𝑛 𝑜𝑓 𝑠𝑙𝑜𝑝𝑒) = 𝑠𝑦𝑥 √∑(𝑥 𝑖 -𝑥 ̂)2 𝑠𝑐 (𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑑𝑒𝑣𝑖𝑎𝑡𝑖𝑜𝑛 𝑜𝑓 𝑖𝑛𝑡𝑒𝑟𝑠𝑒𝑐𝑡𝑖𝑜𝑛) = 𝑠𝑥𝑦 × √ 1 𝑛 + ∑ 𝑥 ̂2 ∑(𝑥 𝑖 -𝑥 ̂)2 𝑠𝑦𝑥 = √ ∑(𝑦 𝑖 -𝑦 𝑐𝑎𝑙𝑐 ) 2 𝑛-2
is the random errors in the y-direction.

The correlation coefficient, so-called r 2 , should be included in the interval 0 ≤ 𝑟 2 ≤ 1, following:

𝑟 2 = 1 - ∑(𝑦 𝑖 -𝑦 𝑐𝑎𝑙𝑐 ) 2 ∑(𝑦 𝑖 -𝑦 ̂)2
The uncertainty of the calculated ycalc was obtained by the application of the standard deviation of the calculated values and the statistical significance of the r 2 of the calibration curve assuming a t-test and 𝑛 -2 degrees of freedom.

The error of ycalc was calculated based on its standard deviation and the probability of its occurrence.

The inferior and superior error of the ycalc follows, respectively:

𝑦 𝑐𝑎𝑙𝑐 𝑖𝑛𝑓 = 𝑦 𝑐𝑎𝑙 -𝑠𝑦 𝑐𝑎𝑙 × P|𝑋| > 𝑡; P|𝑋|~ 2 at 95 % of significance

A.2. TOC Analyser:

The total organic carbon of the samples was measured by a Total Organic Carbon Analyser TOC V-CSH series Shimadzu (Standalone high-sensitivity model). The TOC analyser is an analytical technique to measure the total inorganic and total carbon in liquid (or solid) sample by combustion catalytic oxidation/NDIR (non-dispersive infrared detector) method.

Method of measurement

The total organic carbon is determined indirectly by the difference of the total carbon (TC) and the total inorganic carbon (TIC) measured by the analyser.

𝑇𝑂𝐶 = 𝑇𝐶 -𝑇𝐼𝐶

The TC is measured when a volume of sample is injected through a carrier gas and catalyst-filled TC combustion tube heated to 680°C. The TC in the sample is then oxidized or decomposes to create CO2. After, the combustion product is cooled and dehumidified by a halogen scrubber into the sample cell of the NDIR, where the carbon dioxide is detected.

The TIC is measured by acidifying the sample through a reactor containing 25 % of H3PO4 and transform inorganic carbon in carbon dioxide which is further transported by gas carrier. This carbon dioxide is detected by the NDIR and the TIC is measured in the same way as TC. The TIC is a combination of carbonate and bicarbonate.

Uncertainty of measurement:

The uncertainty of the TOC analysis was calculated considering the measurements of the TC and IC from the ultra-pure water, and their respective standard solutions.

𝑈𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦 (%) = | [𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙] -[𝑚𝑒𝑎𝑠𝑢𝑟𝑒 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑] -[𝑢𝑙𝑡𝑟𝑎𝑝𝑢𝑟𝑒 𝑤𝑎𝑡𝑒𝑟] [𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙] |
The TOC measurements are rather sensitive to the atmospheric conditions, such as humidity and atmospheric pressure. For this reason, the TC and TIC was measured in ultra-pure water and used to correct the TC and TIC measurements of the sample.

A.3. Determination of the alkalinity:

The alkalinity was determined by colorimetric analysis using Bromocresol Green-Methyl Red Indicator Powder Pillow as colour indicator and a Digital Titrator (Hach method 8203) with a digital titrator cartridges containing 0.16 N (or 0.02 N for low concentration of carbonates in the system) of H2SO4.

The measurements of alkalinity were performed in the blank samples in the absence of uranium in solution. The procedure included the following steps:

-Cooling of the sample at room temperature (for all samples carried out at temperature above 20 ºC)

-An aliquot of 10 mL was sampled, filtrated using 0.2 µm cellulose acetate membrane and added to a flask -The colour indicator was added and swirled to mix until all powder was dissolved -The cartridges containing the titrant were hold in a digital titration and turning the digital titrator's knob the titrant is added to the solution. The flask is swirled every time the titrant was added until the colour changed from green to pink.

-The number of digits was recorded from digital titrator counter.

-The alkalinity was determined as following:

𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦 𝑡𝑜𝑡𝑎𝑙 = 𝑁 𝑎𝑐𝑖𝑑 ( 𝑒𝑞 𝐿 ) • 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑑𝑖𝑔𝑖𝑡𝑠 800 (𝑚𝐿) • 1 𝑣𝑜𝑙𝑢𝑚𝑒 𝑠𝑎𝑚𝑝𝑙𝑒 (𝑚𝐿)

A.4. pH measurements

The pH was measured by a pH-meter combined with a pH electrode and thermometer detector (pHC2085, Hach Lange) as hydrogen ion activity. The pH electrode was composed by glass combined with a red rod reference electrode composed by silver/silver chloride (Ag/AgCl). The calibration of the pHmeter was automatically performed by PHM220 Lab pHmeter (Radiometer Analytical) at the beginning of each experiment using 2 standard buffer solutions of pH 4.005 and 7.000 manufactured according to Radiometer Analytical specifications. All standard solutions were stable in a temperature range of 0 to 95ºC. The calibration of the pHmeter were carried out under the temperature of interest using a duo dry-bath device previously heated (Fisher Scientific).

A.5. Stock solutions

All solutions were prepared with ultra-pure deionized water (18.2 MΩ/cm) and with commercially available chemical products of analytical grade or superior.

Calcium, Sodium and carbonate:

The stock solutions of calcium, sodium and carbonate were prepared by dissolving Ca(NO3)2, NaNO3 (or NaClO4, NaCl) and NaHCO3 (Sigma Aldrich, St. Louis MO), respectively, in ultrapure deionised water.

Uranium:

Stock solutions of uranium were prepared for all different batch experiments performed with U(VI). 238 U(VI)-impoverished solution is prepared from ICP-MS standard solution (1000 µg/mL in 4% of HNO3 matrix) (SCP Science, France) or ICP-MS standard solution (10000 µg/mL in 0.4% of HNO3 matrix) (SCP Science, France).

The U(VI) samples, previously acidified to pH 2 with HNO3 ultrapure (70%) were measured by ICP-MS (ThermoScientific XSERIES 2 ICP-MS).

Synthetic Callovo-Oxfordian porewater:

The synthetic Callovo-Oxfordian porewater was prepared according to the reported data from BRGM (2008). The chemical composition of the synthetic porewater follows the reduced mineralogical system at 20 and 80ºC. The chemical composition was corrected by previous calculations using PhreeqC by equilibrating the solution in pCO2 of 10 -2.1 and 10 -1.3 at 20 and 80 ºC, respectively (Table A. 2). Therefore, stock solutions were prepared in controlled atmosphere of 1% or 5% of CO2(g) (accordingly to the pH/pCO2/T) at room temperature by dissolving in 1L of degassed ultrapure deionised water the salts described in The gel-type Dowex 1x2 strong base anion exchange resin was purchased from Sigma-Aldrich (France) in chloride form. This resin is characterised by a total exchange capacity of 0.7 meq/mL by wetted bed volume; 70 %wt water retention capacity; and density of 44 lbs/ft 3 ≈ 0.705 g/cm 3 . Primarily, the resin is converted to nitrate form by equilibrating with 250 mL of 1 M NaNO3 (or NaClO4) during 1 day on a magnetic shaker. The supernatant is filtrated and replaced by 250 mL of NaNO3 (or NaClO4) (1 M) and the process was repeated for 3 times. After 3 equilibration cycles, the resin is filtrated and rinsed with deionised water after overnight equilibration. After rising overnight, the resin is once again filtrated and air dried.

Callovo-Oxfordian claystone:

Calovo-Oxfordian claystone samples were extracted from the borehole EST51779 drilled in the Callovo-Oxfordian formation. The sample was selected and provided by Andra in July 2015. It has always been storage in closed and clean atmosphere to avoid any oxidation and carbonatation of the mineral surface. The composition is provided in Table A. 4.

Illite:

Illite is originated from Oligocene geological formation from Le Puy-en-Valey (Haute-Loire, France). The so-called illite du Puy was purified and provided by BRGM. The purification procedure was also performed by BRGM.

A.7. Batch experiments procedures

Batch experiments are performed to: 1) determine the stability constants of Ca-U(VI)-CO3 at different temperatures; 2) to study the main mechanisms of retention of uranium onto Callovo-Oxfordian formation; and 3) the respective effect of the temperature on these retention mechanisms. The batch experiments are performed in both open and closed atmosphere. The experimental procedure in open or closed system was applied in the same way for all batch. The details of the experimental conditions for each batch experiment are described in Table A. 5. The procedure at open system consisted in firstly weighting the solid phase in 35 mL PPCO tubes at room temperature and prepare the stock solutions with ultrapure water.

Anionic-exchange batch experiments were performed in atmospheric conditions to test the efficiency of using NaClO4 instead of NaNO3. In this case, the initial pH controlled by adding 0.1 M of NaHCO3 and was adjusted for each tube by adding 0.01 M of HNO3 or NaOH until the pH 8.1 was reached.

Concerning the sorption edge batch experiments, the pH of the stock solution of 0.1 NaCl was controlled by bubbling with Na2(g)/CO2(g) (99/1 ratio) the solution until reach pH of interest. Afterwards, the solutions were agitated over night to favour the equilibrium with atmospheric pCO2. In each tube, the pH of the solution was measured and adjusted by adding 0.01 M of HCl or NaOH. The results of the pH of each stock solution are shown in Table C. 8.

Each tube was sealed with Teflon-film to avoid any leaking. The samples were transferred to the incubator coupled with a reciprocating shaker (Thermotron InforsHT) and let in equilibrium for 24 h by agitation at temperature of interest. After 24 h, each sample was introduced in the thermo-Teflonaluminium box, specially prepared to maintain the temperature of the samples, and an aliquot of a solution of uranium was added. Each tube was closed and transferred once more to the incubator at the temperature of interested. After 3 days of equilibration, the samples were introduced again in the thermo-Teflon-aluminium box, and an aliquot of 5 mL was sampled by 5 mL syringe PP and filtrated by 0.22 μm filters of cellulose acetate (20 ºC) or Millex IC (higher temperatures), discarding the first 1-2 mL, acidified to pH 2 with HNO3 ultrapure (70%) and saved for further analysis of uranium. The final pH was measured in the remaining solution just after the sampling to avoid the loss of heat.

Two blanks were prepared with and without solid or uranium in solution. The alkalinity was measured at end of the experiments when the temperature of the sample was cooled down using blank sample without U in solution.

A.7.2. Closed system (controlled atmosphere) procedure:

In the experiments performed in closed system, the samples were weight in room conditions in 35 mL PPCO tubes and transferred into the glove-box Ar(g) or Ar(g)/CO2(g) (with CO2(g) percentage according to temperature of interest). Simultaneously, 1 L of ultrapure water was degassed by agitating and heating at approx. 50 ºC using a vacuum pumping system and transferred into the glove-box. The stock solutions were prepared individually using degassed ultrapure water in the glove-box and added to each tube.

In case of the sorption edge, the pH of the solutions was adjusted for each tube by adding different aliquots of 0.01 M of HNO3 or NaOH (Table C. 38 and Table C. 39).

For the anionic-exchange experiments and sorption isotherms, the pH was controlled by adding 0.1M of NaHCO3.

Each tube was sealed with Teflon-film to avoid any leaking, closed and removed from the glove-box. The samples were quickly packed in vacuum by wrapping with an aluminium-envelope using a Vacupac system. After, the packed samples were transferred into an incubator coupled with reciprocating shaker (Thermotron InforsHT) heated previously to the temperature of interested and left in equilibrium for 1 day.

After 24 h of equilibration, the aluminium-envelopes were opened, and each sample was introduced inside of a thermo-Teflon-aluminium box, specially prepared to maintain the temperature of the samples, and transfer into the glove-box as quick as possible. In the glove-box an aliquot of a solution of U(VI). Each tube was closed and removed from the glove-box, and quickly packed in an aluminiumenvelope once more in vacuum. The packed samples were transferred into the incubator coupled with reciprocating shaker and left in equilibration.

After equilibration, the aluminium-envelope was opened, and each sample was introduced inside of a thermo-Teflon-aluminium box and transferred into the glove-box, and 5 mL were sampled by 5 mL syringe PP and filtrated by 0.22 μm filters of cellulose acetate (20 ºC) or Millex IC (higher temperatures), discarding the first 1-2 mL, acidified to pH 2 with HNO3 ultrapure (70%) and saved for further analysis of uranium. The final pH was measured in the remaining solution just after the sampling to avoid the loss of heat. Two blanks were prepared with and without solid or uranium in solution. The alkalinity was measured at end of the experiments when the temperature of the sample was cool down to room temperature using blank sample without U in solution (see Table C. 9, Table C. 23, Table C. 44 and Table C. 51).

Assuming that the fraction of illite and smectite in the interstratified illite/smectite corresponds each to 50%, then:

𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑒𝑥𝑐ℎ𝑎𝑛𝑔𝑒 𝑠𝑖𝑡𝑒𝑠 (𝑚𝑜𝑙) = 𝑤𝑡 𝑐𝑙𝑎𝑦 % 100 • 0.5 • 𝐶𝐸𝐶 ( 𝑒𝑞 𝑘𝑔 ) • 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒(𝑘𝑔)
Where, CEC corresponds to the cationic exchange capacity.

B.3. Calculation of the concentration of U(VI) sorbed on the solid:

The concentration of U(VI) sorbed on the solid was calculated based on ICP-MS measurements and follows:

B.4. Calculation of distribution coefficient of U(VI):

The distribution coefficient of U(VI) is calculated as follow: The pH measurements were performed in the glove-box, except for the test 0, which was measured in atmospheric conditions. The results show that for the pH measured in controlled atmosphere no changes were observed, which is constant at pH 7.3 (± 0.03). However, for the pH measured in atmospheric conditions the pH change with time (Figure C. 5).

As conclusion, the packing of the tubes using aluminised envelops (under vacuum or CO2(g) atmosphere) is necessary and is effective as the glove-box, and the pH must be measured under a controlled atmosphere to avoid any changes of chemical equilibrium. Total Organic Carbon:

The tubes of PPCO may release organic carbon due to deterioration by increasing the temperature. This organic carbon might influence the speciation of U(VI) in the chemical system. For this reason, some tests were performed to ensure the minor impact of the organic carbon released by the degradation of the tubes.

The experimental tests were prepared according the time of washing and the number of washes need to release the maximum of organic carbon from the tubes. An additional test was performed using 5% of HNO3 to rinse the tubes before starting the washes process.

In the beginning of each analysis, the uncertainty of the signal is calculated considering the TC and IC from ultra-pure water and the respective standard solutions. The uncertainty of the signal of the TC and IC are expressed in the Table C. 5. All analytical results are summarised in Table C. 6.

The number of washes was limited to TOC < 1 ppm. At 20 ºC, this limit is reached after one wash, however, in temperature, it was only after 3 washes (Figure C. 7a). For all temperatures, the three washes procedure can be done in 1 day, no significant difference is observed increasing the timing between washes (Figure C. 7).

The results also show that use the acidic rinse before the washes does not affect the release of the organic carbon from the PPCO tubes in comparison with rinsing with ultra-pure water before washing process (Figure C. 8b). As a conclusion of these experimental tests, a cycle of 3 washes with ultrapure water was carried out at the temperature of interest before using the PPCO tubes in the batch experiments. A bottom-up approach with the published "2SPNE SC/CE" model was used for describing the sorption processes, with the assumption that the clay fraction of the COx (Illite, and I/S) governs U(VI) sorption.

The model was successfully applied to reproduce a wealth of experimental data obtained with illite, the COx clay fraction and the COx clay rock as a function of key parameters (pH, pCO2, [U(VI)], [Ca]) at 20 °C. The sorption on COx conditions is mainly governed by the sorption of U(VI)-CO3 complexes and a new sorption constant is proposed for illite.

An increase in temperature to 80 °C leads to an in-crease of U(VI) retention on COx. This increase is ac-companied by a change of both pCO2 and pH. The sorption model developed at 20 °C, together with the thermodynamic parameters describing U(VI) speciation in solution, can explain this increase but without obtaining a good agreement with the experiment. The model is improved by considering ΔrHº values for sur-face complexation reactions obtained for the U(VI))/illite system.

Figure 1 .Figure 6 .Figure 7 .Figure 8 .Figure 9 .

 16789 Figure 1. Schematic representation of deep geological nuclear waste disposal multi-barrier system (1) Callovo-Oxfordian host-rock; (2) engineered barrier for IWL; (3) engineered barrier for HWL; (4) and (5) HWL primary package (e.g. vitrified waste). ............................................................................................................................. 30 Figure 2. Schematic pathway for the safety assessment of the nuclear disposal. ..................................................... 31 Figure 3. Schematic representation of U(VI) controlling processes in radioactive waste disposal system using hostclay rock. The main processes concerning in the present study are highlighted in red. ..................................... 32 Figure 4. Geological location of the underground research laboratory within the Paris Basin (upper) and cross section of the geological formations surrounding the underground research laboratory. Adapted from Linard et al., 2011. ............................................................................................................................................................. 33 Figure 5. Calculated U(VI) speciation as a function of calcium concentrations between 0 and 5•10 -3 M at 25 ºC, [U(VI)]total = 1•10 -8 M, [CO3 2-]total = 3•10 -3 M, pH = 7.2 and ionic strength of 0.1 M using NaNO3 as electrolyte. Calculations performed with ThermoChimie database v9b. ............................................................................... 41 Figure 6. Scheme of the experimental set-up for the closed system experiments. ................................................... 47 Figure 7. Experimental data originated from supporting information of Dong and Brooks (2006). Fit approach of the a) U(VI) sorbed onto resin as function of U(VI) in equilibrium in solution and b) the distribution of U(VI) as a function of calcium. ..................................................................................................................................... 50 Figure 8. Mineral saturation index as a function of Ca considered for a) Dong and Brooks (2006) using ThermoChimie database Davies (2009) and b) Endrizzi and Rao (2014) using ThermoChimie database SIT (2014). ................................................................................................................................................................. 52 Figure 9. Concentration of U(VI) in solution as a function of the concentration of calcium in solution at pH of 7.8 (± 0.1), total [U(VI)] = 4.7 (± 0.3)•10 -7 M, I = 0.1 (± 0.005) M NaClO4, in atmospheric conditions and 20 °C. ............................................................................................................................................................................ 54 Figure 10. Measured pH as a function of S/L ratio in the experiments in the absence of calcium at a) 20 ºC in equilibrium with 1% of CO2(g) and b) 80 ºC in equilibrium with 5% of CO2(g). .............................................. 55 Figure 11. Speciation distribution of U(VI) in the absence of calcium in solution as a function of the temperature. ............................................................................................................................................................................ 56 Figure 12. Concentration of U(VI) in the absence of calcium in solution as a function of the solid-liquid ratio (g/L) at a) 20 ºC in equilibrium with 1% of CO2(g); b) 40 ºC in equilibrium with 1% of CO2(g); c) 60 ºC in equilibrium with 3% of CO2(g); and d) 80 ºC in equilibrium with 5% of CO2(g). Solid-line: Model fitting of the concentration of U(VI) as a function of the solid-liquid ratio at 20, 40, 60 and 80ºC. .............................................................. 55 Figure 13. Dependency of the calculated log10KX1 value from Reaction 10 with the temperature. ......................... 57 Figure 14. Concentration of U(VI) in equilibrium in solution as a function of calcium at different temperatures in 0.1 M of NaNO3. ................................................................................................................................................. 58 Figure 15. Symbols: Measured concentration of U(VI) for the Type II experiments as a function of the concentration of calcium at a) 20 ºC and b) 40 ºC in equilibrium with 1% of CO2(g); c) 60 ºC in equilibrium with 3% of CO2(g); and d) 80 ºC in equilibrium with 5% of CO2(g). Solid-line: calculated U(VI) concentration using the model proposed in the present work. ............................................................................................................................. 59

Figure 32

 32 Figure 32. a) Callovo-Oxfordian claystone sample: Sorption of 6.8 (±0.2)•10 -7 M of U(VI) as a function of time in 1% of CO2(g), synthetic COx porewater, pH of 7.3 (± 0.1) and S/L of 24 g/L. b) Illite: Sorption of 9.8 (±0.8)•10 - 8 M of U(VI) as a function of time in atmospheric pCO2, 0.1 M of NaCl, pH 6.3 (± 0.1) and S/L of 0.5 g/L. .. 86 Figure 33. Speciation distribution of the aqueous uranyl species (left) and uranyl hydrolysis surface complexes (right) as function of the pH in the absence of carbonates at 0.1 M of NaNO3 in contact with illite with S/L of 2 g/L and U(VI)initial of 1.1(±0.05)•10 -7 M. ............................................................................................................ 88 Figure 34. Comparison of the sorption edges of U(VI) on Na-illite in the absence of carbonates between p.w.: present work (0.1 M of NaNO3, S/L of 2 g/L and U(VI)initial of 1.1(±0.05)•10 -7 M) and [1]Bradbury and Baeyens (2009) (0.1 M of NaClO4, S/L 2.6 g/L and U(VI)initial of 10 -7 M). ................................................................................. 88 Figure 35. Sorption isotherm of U(VI) onto Na-illite in the absence of carbonates in Ar(g) atmosphere at 0.1 M of NaNO3, S/L of 2 g/L and pH of 6.4. ................................................................................................................... 89Figure 36. Sorption isotherms of U(VI) on Na-illite in the absence of carbonates between present work (0.1 M of NaNO3 in S/L of 2 g/L and pH of 6.4) and work of Bradbury and Baeyens (2005) (0.1 M of NaClO4, S/L 2.6 g/L and pH of 5.8). .............................................................................................................................................. 89 Figure 37. Sorption edges of U(VI) on illite in 0.1 M of NaNO3, S/L of 2 g/L and U(VI)initial of 1.1(±0.1)•10 -7 M, in the absence (open circles) and presence (blue circles) of carbonates in equilibrium with atmospheric CO2. .... 90 Figure 38. Speciation distribution of the aqueous uranyl species (left) and uranyl hydrolysis and carbonato surface complexes (right) as function of the pH in the absence of carbonates at 0.1 M. ................................................ 91 Figure 39. Model approach on sorption edge of U(VI) on illite. Model 0: neither ≡SsOUO2CO323or ≡Ss,wOUO2CO3are considered; Model I: considering the logK from Montmorillonite; Model II: only ≡SsOUO2CO323with log10 K = 17.5 (±0.5) is considered; and Model III: the ≡SsOUO2CO323-, the ≡SsOUO2CO3and ≡SwOUO2CO3are considered, with log10 K of 17.5, 9.8 and 9.3, respectively. Error bars that are not visible are smaller than the symbol size. .......................................................................................... 92 Figure 40. Sorption isotherms of U(VI) onto illite in the absence of carbonates (o) at equilibrium pH of 6.4 in 0.1 M of NaCl and presence of carbonates in equilibrium with atmospheric pCO2 (•) at equilibrium pH of 7.3 in 0.1 M of NaNO3. Error bars that are not visible are smaller than the symbol size. .................................................. 93 Figure 41. Prediction model of sorption of U(VI) onto illite in equilibrium with pCO2 atmospheric in 0.1 M of NaNO3 at pH 7.3. a) Model II: only ≡SsOUO2CO323in considered; b) Model III: U(VI)-CO3 surface complexation reactions from Table 24. Error bars that are not visible are smaller than the symbol size. .......... 93 Figure 42. Sorption isotherms of U(VI) onto illite in the presence of carbonates in equilibrium with atmospheric pCO2, 1% and 5% of CO2(g) in 0.1 NaNO3, pH fixed at 7.3 (• and •), and 6.6 (•), respectively. Error bars that are not visible are smaller than the symbol size. ................................................................................................. 94 Figure 43. Model approach for the sorption isotherms of U(VI) onto illite in 0.1 M of NaNO3, S/L of 2 g/L at pH 7.3 in equilibrium with 1% of CO2(g) (left top and bottom) and pH 6.6 in equilibrium with 5% of CO2(g) (right top and bottom). Model II: only ≡SsOUO2CO323in considered; Model III: U(VI)-CO3 surface complexation reactions from Table 24. Error bars that are not visible are smaller than the symbol size. ................................ 95 Figure 44. Sorption isotherms of U(VI) on illite in the presence of 1% of CO2(g) under 0.1 M NaNO3 at pH 7.3 and synthetic COx porewater at pH 7.1. Error bars that are not visible are smaller than the symbol size. ............... 96 Figure 45. Sorption isotherms of U(VI) on illite as a function of the U(VI) in COx porewater in equilibrium with 1% of CO2(g). Model II: only ≡ 𝑺𝒔𝑶𝑼𝑶𝟐𝑪𝑶𝟑𝟐𝟑in considered taking also in account the thermodynamic data for Ca-U(VI)-CO3 ternary complexes (see Chapter 2). Error bars that are not visible are smaller than the symbol size. ........................................................................................................................................................ 96 Figure 46. a) Sorption isotherm of U(VI) onto (o) illite and (◊) Callovo-Oxfordian clay fraction (illite + I/S) in the absence of carbonates in both 2 g/L under 0.1 M NaNO3 at pH of 6.4 and 7.0 (respectively). b) Sorption isotherm of U(VI) onto (•) illite and (♦) Callovo-Oxfordian clay fraction (illite + I/S) in the presence of carbonates in equilibrium with 1% of CO2(g) in both 2 g/L under 0.1 M NaNO3 at pH of 7.3 and 7.2 (respectively). Model approach for illite (solid-line) and for COx clay fraction (dash-line). Error bars that are not visible are smaller than the symbol size. ........................................................................................................................................... 98

  Figure 59. Model approach on sorption of isotherms of U(VI) in the absence of carbonates at a) 20 ºC; b) 60 ºC; and c) 80 ºC. Model II: model approach without enthalpy values (see Chapter 3); Model IV: model approach considering the enthalpy values (see Table 33). Error bars that are not visible are smaller than the symbol size. .......................................................................................................................................................................... 119 Figure 60. Distribution of U(VI) hydrolysis surface complexes of illite at 20ºC (solid-lines) and 80ºC (dashed-lines) as a function of the pH. Calculations taking the parameters given in Table 33. ............................................... 120 Figure 61. Fitting model for sorption edge experiments (a) at 20 ºC and (b) at 80 ºC (with ±0.2 log units of uncertainty for the stability constants); for sorption isotherm experiments (c) at 20 ºC (with ±0.2 log units of uncertainty for the stability constants); (d) at 60 ºC and (e) at 80 ºC (with ±0.5 log units of uncertainty for the stability constants). Error bars that are not visible are smaller than the symbol size. ....................................... 121

Figure 62 .

 62 Figure62. Relationship between Log10Kºx of aqueous species and Log10 S Kºx of surface complexes for illite determined at 20, 60 and 80ºC in comparison with the respective standard stability constants at 25ºC (solid-line) for illite reported byBradbury and Baeyens (2009). ........................................................................................ 122 Figure 63. Relation of constant of hydrolysis surface complexation (T) and standard constant of hydrolysis surface complexation (T0) as a function of the 1/T0-1/T for the surface species a) ≡SsOUO2+, b) ≡SsOUO2OH, c) ≡SsOUO2(OH)2-, and d) ≡SsOUO2(OH)32-. Symbols: experimental data; Dashed-lines: regression line calculated for the enthalpy values. .................................................................................................................... 122 Figure 64. Symbols: measured experimental data from sorption edge and isotherm of U(VI) at different temperatures. Solid-lines: calculated U(VI) concentration using enthalpy values calculated from Table 35. Dashed-lines correspond to the uncertainty of the enthalpy values. Error bars that are not visible are smaller than the symbol size. ................................................................................................................................................. 123 Figure 65. Distribution coefficient of U(VI) as a function of the pH in 0.1 NaCl a) in the absence or presence of carbonates in equilibrium with pCO2 atmospheric at 80 ºC. Error bars that are not visible are smaller than the symbol size. ...................................................................................................................................................... 124 Figure 66. Sorption isotherms of U(VI) on illite at different temperatures in equilibrium with a) Ar(g); b) pCO2 atmospheric; c) pCO2 of 1% of CO2(g); and d) pCO2 of 5% of CO2(g). Error bars that are not visible are smaller than the symbol size. ......................................................................................................................................... 125 Figure 67. Distribution of the speciation U(VI) in solution as a function of the pH at 20 (solid-lines) and 80 ºC (dashed-lined) in equilibrium with pCO2 atmospheric. .................................................................................... 126 Figure 68. Sorption edge of U(VI) in equilibrium with pCO2 atmospheric. Model V: enthalpy values considered only for hydrolysis surface complexation reactions (Table 35). ....................................................................... 127 Figure 69. Model approach considering only the enthalpy of the hydrolysis of surface complexes at different temperatures in the presence of carbonates (see Table 35). Dashed-lines correspond to the uncertainty of the enthalpy values of the hydrolysis surface complexes. ...................................................................................... 127 Figure 70. Sorption isotherms of U(VI) on illite and Callovo-Oxfordian clay fraction (illite + I/S) system at 80 ºC a) in the absence of carbonates and b) in the presence of carbonates in equilibrium with pCO2 of 5% of CO2(g). .......................................................................................................................................................................... 128 Figure 71. Model approach for describing the sorption isotherms of U(VI) on Callovo-Oxfordian clay fraction (illite + I/S) at 20 and 80 ºC a) in the absence and b) in the presence of carbonates in equilibrium with 1% and 5% of CO2(g), respectively. Taking the thermodynamic parameters for illite and montmorillonite (see Table 39) the solid-lines correspond to the model prediction at 20ºC; and the dashed-lines correspond to the model prediction at 80 ºC. ............................................................................................................................................................. 129 Figure 72. Sorption isotherms of the U(VI) in synthetic porewater performed at 20 and 80 ºC for a) U(VI)/illite system; b) U(VI)/ Callovo-Oxfordian clay fraction system; and c) U(VI)/ Callovo-Oxfordian claystone system.Model approach considering thermodynamic data for surface reactions from Table39and using the thermodynamic data for Ca-U(VI)-CO3 aqueous complexes (see Table14). Solid-line: model at 20 ºC; dashedline: model at 80ºC. ........................................................................................................................................... 131

  Figure 72. Sorption isotherms of the U(VI) in synthetic porewater performed at 20 and 80 ºC for a) U(VI)/illite system; b) U(VI)/ Callovo-Oxfordian clay fraction system; and c) U(VI)/ Callovo-Oxfordian claystone system.Model approach considering thermodynamic data for surface reactions from Table39and using the thermodynamic data for Ca-U(VI)-CO3 aqueous complexes (see Table14). Solid-line: model at 20 ºC; dashedline: model at 80ºC. ........................................................................................................................................... 131 Figure 73. Impact of the temperature on the distribution coefficient of U(VI) as a function of the concentration of U(VI) in equilibrium in Callovo-Oxfordian claystone system. (•) and (□) experimental data from the present work. red-line: model approach at 80ºC considering the model approach proposed in the present study and the thermodynamic parameters of U(VI) aqueous speciation (Table 14). .............................................................. 131

Figure 1 .

 1 Figure 1. Schematic representation of deep geological nuclear waste disposal multi-barrier system (1) Callovo-Oxfordian host-rock; (2) engineered barrier for IWL; (3) engineered barrier for HWL; (4) and (5) HWL primary package (e.g. vitrified waste).

Figure 2 .

 2 Figure 2. Schematic pathway for the safety assessment of the nuclear disposal.

Figure 3 .

 3 Figure 3. Schematic representation of U(VI) controlling processes in radioactive waste disposal system using host-clay rock. The main processes concerning in the present study are highlighted in red.

Figure 4 .

 4 Figure 4. Geological location of the underground research laboratory within the Paris Basin (upper) and cross section of the geological formations surrounding the underground research laboratory. Adapted from Linard et al., 2011.

  in order to explain U(VI) sorption, mobility and speciation in different systems. For example,[START_REF] Meleshyn | Influence of (Calcium-)Uranyl-Carbonate Complexation on U(VI) Sorption on Ca-and Na-Bentonites[END_REF] evaluated the influence of Ca-U(VI)-CO3 complexes on uranium sorption on bentonites;[START_REF] Prat | Uranium Speciation in Drinking Water from Drilled Wells in Southern Finland and Its Potential Links to Health Effects[END_REF] included Ca-U(VI)-CO3 complexes in their studies for U(VI) speciation in drinking water; and[START_REF] Stewart | Impact of Uranyl-Calcium-Carbonato Complexes on Uranium(VI) Adsorption to Synthetic and Natural Sediments[END_REF] studied the impact of Ca-U(VI)-CO3 complexes on uranium adsorption to synthetic and natural sediments.The Ca-U(VI)-CO3 ternary complexes and their associated stability constants (as reported in 2006 by Dong and Brooks) have already been included in the ThermoChimie database version 9.b(Giffaut et al., 2014, www.thermochimie-tdb.com). Using this thermodynamic database, we were able to calculate the distribution of the aqueous speciation of uranium under specific geochemical conditions (Figure5). From these results, it is possible to evaluate the importance of such aqueous species on the behaviour of U(VI) under neutral pH and in the presence of calcium and carbonates systems as clay systems, relevant for nuclear waste disposal.

Figure

  Figure 5. Calculated U(VI) speciation as a function of calcium concentrations between 0 and 5•10 -3 M at 25 ºC, [U(VI)]total = 1•10 -8 M, [CO3 2-]total = 3•10 -3 M, pH = 7.2 and ionic strength of 0.1 M using NaNO3 as electrolyte. Calculations performed with ThermoChimie database v9b.

Figure 6 .

 6 Figure 6. Scheme of the experimental set-up for the closed system experiments.

Figure 7 .

 7 Figure 7. Experimental data originated from supporting information of Dong and Brooks (2006). Fit approach of the a) U(VI) sorbed onto resin as function of U(VI) in equilibrium in solution and b) the distribution of U(VI) as a function of calcium.

Table 9 .

 9 Theoretical and measured parameters from the evaluation of CO2(g) in equilibrium at different temperatures. Theor.: theoretical value in agreement with the initial conditions of the experiments; Exp.: pH and alkanity measured at the end of the experiment, and pCO2 calculated according pH and alkanity measurements.

Figure 8 .

 8 Figure 8. Mineral saturation index as a function of Ca considered for a) Dong and Brooks (2006) using Thermo-Chimie database Davies (2009) and b) Endrizzi and Rao (2014) using ThermoChimie database SIT (2014).

Figure 9 .

 9 Figure 9. Concentration of U(VI) in solution as a function of the concentration of calcium in solution at pH of 7.8 (± 0.1), total [U(VI)] = 4.7 (± 0.3)•10 -7 M, I = 0.1 (± 0.005) M NaClO4, in atmospheric conditions and 20 °C.

Figure 10 .

 10 Figure 10. Measured pH as a function of S/L ratio in the experiments in the absence of calcium at a) 20 ºC in equilibrium with 1% of CO2(g) and b) 80 ºC in equilibrium with 5% of CO2(g).

Figure 11 .

 11 Figure 11. Concentration of U(VI) in the absence of calcium in solution as a function of the solid-liquid ratio (g/L) at a) 20 ºC in equilibrium with 1% of CO2(g); b) 40 ºC in equilibrium with 1% of CO2(g); c) 60 ºC in equilibrium with 3% of CO2(g); and d) 80 ºC in equilibrium with 5% of CO2(g). Solid-line: Model fitting of the concentration of U(VI) as a function of the solid-liquid ratio at 20, 40, 60 and 80ºC.

Figure 12 .

 12 Figure 12. Speciation distribution of U(VI) in the absence of calcium in solution as a function of the temperature.

Figure 13 .

 13 Figure 13. Dependency of the calculated log10KX1 value from Reaction 11 with the temperature.

Figure 14 .

 14 Figure 14. Concentration of U(VI) in equilibrium in solution as a function of calcium at different temperatures in 0.1 M of NaNO3.

Figure 15 .

 15 Figure 15. Symbols: Measured concentration of U(VI) for the Type II experiments as a function of the concentration of calcium at a) 20 ºC and b) 40 ºC in equilibrium with 1% of CO2(g); c) 60 ºC in equilibrium with 3% of CO2(g); and d) 80 ºC in equilibrium with 5% of CO2(g). Solid-line: calculated U(VI) concentration using the model proposed in the present work.

Figure 16 .

 16 Figure 16. Symbols: Experimental KD of U(VI) for the experiments Type II as function of the concentration of calcium at a) 20 ºC and b) 40 ºC in equilibrium with 1% of CO2(g); c) 60 ºC in equilibrium with 3% of CO2(g); and d) 80 ºC in equilibrium with 5% of CO2(g). Solid-line: calculated U(VI) concentration using the model proposed in the present work.

Figure 17 .

 17 Figure 17. Temperature effect of the formation of CaUO2(CO3)3 2-(left) and Ca2UO2(CO3)3(aq) (right).

Figure 18 .

 18 Figure 18. 2D mapping of the stability constants of Ca-U(VI)-CO3 aqueous complexes distribution at 20, 40, 60 and 80 ºC under respective pCO2. Squares: optimised values using PEST+Phreeqc; Circles: optimised values using Excel+Phreeqc.

Figure 19 .

 19 Figure 19. Dependency of the calculated log10β value for Reaction 1 (left) and Reaction 2 (right) with the temperature using PEST-PhreeqC (black circles) and Excel-PhreeqC (white squares).

in

  All stability constants determined by PEST or 2D mapping are calculated at infinite dilution (zero ionic strength) and corrected using the Davies equation for each temperature stated; values are plotted as Figure20. PEST and 2D mapping results are in excellent agreement in the case of CaUO2(CO3)3 2-although the stability constants for Ca2UO2(CO3)3(aq) have a higher uncertainty (Figure20).

Figure 20 .

 20 Figure 20. Variation of CaUO2(CO3)3 2-(left) and Ca2UO2(CO3)3(aq) (right) stability constants with the temperature. Black: PEST approach. Red: 2D mapping approach. The lines represent least squares fitting following the constant enthalpy of the reaction equation.

Figure 21 .

 21 Figure 21. Symbols: measured U(VI) concentration as a function of the concentration of calcium in solution for Type II experiments at 20, 40, 60 and 80 ºC. Solid line: Calculated uranium concentration using the standard stability constant and enthalpy values from Table14.

Figure 23 .

 23 Figure 23. Calculated uranium aqueous speciation of the concentration of calcium in solution for Type II experiments at 20, 40, 60 and 80 ºC.

Figure 24 .

 24 Figure 24. Comparison of the stability constants at 25 ºC determined from different studies with the stability constants determined in the present work (p.w.). KAL00: Kalmykov and Choppin (2000); BER01: Bernhard et al. (2001); DON06: Dong and Brooks (2006); LEE13: Lee and Yun (2013); END14: Endrizzi and Rao (2014); JO17: Jo and Yun (2017); p.w.: present work.

Figure 25 .

 25 Figure 25. Calculated U(VI) aqueous speciation at calcium concentrations between 0 and 5•10 -3 M at 25ºC. [U]T = 1•10 -8 M, [C]T = 3•10 -3 M, pH = 7.2 and I = 0.1 M. Left: Calculations performed using the stability constants from Dong and Brooks for Ca-U(VI)-CO3 complexes. Right: Calculations performed using the stability constants determined in the present work for Ca-U(VI)-CO3 complexes.

Figure 26 .

 26 Figure 26. Symbols: measured U(VI) concentration as a function of the concentration of calcium in solution for experiments at 20, 40, 60 and 80ºC performed in present work. Blue-line: Calculated uranium concentration using the thermodynamic data for Ca-U(VI)-CO3 complexes determined in present work. Red-line: calculated uranium concentration using the thermodynamic data for Ca-U(VI)-CO3 complexes reported in [1] Endrizzi and Rao (2014).

Figure 27 .

 27 Figure 27. Representation of a) aluminium oxyde octahedra, b) silicium oxyde tetrahedra, c) magnesium oxyde octahedra, plan formed by combination of one tetrahedral sheet with d) one dioctrahedral sheet, and e) one trioctahedrical sheet. (Adapted from Ngouana, 2014).

Figure 28 .

 28 Figure 28. Structural elements of phyllosilicates of tedrahedral and octahedral sheets linked to form a 1:1 layer type structure and linked to form 2:1 layer type structure (Adapted from Sposito et al., 1999; Copyright (1999) National Academy of Sciences http://www.pnas.org/content/96/7/3358).

Figure 29 .

 29 Figure 29. Regular and random interstratified clay minerals. A and B are layers with different periodicity along the c direction (Brigatti et al., 2006).

Figure 30 .

 30 Figure 30. Schematic representation of cation exchange with one uranyl cation and two Na+ cations, and uranyl surface complexation bond to the edge of surface.

Figure 31 .

 31 Figure 31. Schematic representation of the placements of ions, charges (σ) potential (Ψ) and capacitances (C) for different SMCs for surface 0-plane; β-plane and d-plane. Adapted from Goldberg (2013).

Figure

  Figure 32. a) Callovo-Oxfordian claystone sample: Sorption of 6.8 (±0.2)•10 -7 M of U(VI) as a function of time in 1% of CO2(g), synthetic COx porewater, pH of 7.3 (± 0.1) and S/L of 24 g/L. b) Illite: Sorption of 9.8 (±0.8)•10 -8 M of U(VI) as a function of time in atmospheric pCO2, 0.1 M of NaCl, pH 6.3 (± 0.1) and S/L of 0.5 g/L.

Figure 33 .

 33 Figure 33. Speciation distribution of the aqueous uranyl species (left) and uranyl hydrolysis surface complexes (right) as function of the pH in the absence of carbonates at 0.1 M of NaNO3 in contact with illite with S/L of 2 g/L and U(VI)initial of 1.1(±0.05)•10 -7 M.

Figure 34 .

 34 Figure 34. Comparison of the sorption edges of U(VI) on Na-illite in the absence of carbonates between present work (p.w.): present work (0.1 M of NaNO3, S/L of 2 g/L and U(VI)initial of 1.1(±0.05)•10 -7 M) and [1] Bradbury and Baeyens (2009) (0.1 M of NaClO4, S/L 2.6 g/L and U(VI)initial of 10 -7 M).The results of the sorption isotherms as a function of U(VI) concentration at constant pH are shown in Figure35. The sorption of U(VI) onto illite in the absence of carbonates follows a linear isotherm with an increase of U(VI) sorption when U(VI) increases in solution. The measured pH of equilibrium was 6.4 and the alkalinity was 7•10 -5 eq/L (indicating negligible contamination of CO2(g)).

Figure 35 .Figure 36 .

 3536 Figure 35. Sorption isotherm of U(VI) onto Na-illite in the absence of carbonates in Ar(g) atmosphere at 0.1 M of NaNO3, S/L of 2 g/L and pH of 6.4.

Figure 37 .

 37 Figure 37. Sorption edges of U(VI) on illite in 0.1 M of NaNO3, S/L of 2 g/L and U(VI)initial of 1.1(±0.1)•10 -7 M, in the absence (open circles) and presence (blue circles) of carbonates in equilibrium with atmospheric CO2.

Figure 38 .

 38 Figure 38. Speciation distribution of the aqueous uranyl species (left) and uranyl hydrolysis and carbonato surface complexes (right) as function of the pH in the absence of carbonates at 0.1 M.

Figure 39 .

 39 Figure 39. Model approach on sorption edge of U(VI) on illite. Model 0: neither ≡S s OUO 2 (CO 3 ) 2 3-or ≡S s,w OUO 2 CO 3 are considered; Model I: considering the logK from Montmorillonite; Model II: only ≡S s OUO 2 (CO 3 ) 2 3-with log10 K = 17.5 (±0.5) is considered; and Model III: the ≡S s OUO 2 (CO 3 ) 2 3-, the ≡S s OUO 2 CO 3 -and ≡S w OUO 2 CO 3 -are considered, with log10 K of 17.5, 9.8 and 9.3, respectively. Error bars that are not visible are smaller than the symbol size.

Figure 40 .

 40 Figure 40. Sorption isotherms of U(VI) onto illite in the absence of carbonates (o) at equilibrium pH of 6.4 in 0.1 M of NaCl and presence of carbonates in equilibrium with atmospheric pCO2 (•) at equilibrium pH of 7.3 in 0.1 M of NaNO3. Error bars that are not visible are smaller than the symbol size.

Figure 41 .

 41 Figure 41. Prediction model of sorption of U(VI) onto illite in equilibrium with pCO2 atmospheric in 0.1 M of NaNO3 at pH 7.3. a) Model II: only ≡S s OUO 2 (CO 3 ) 2 3-in considered; b) Model III: U(VI)-CO3 surface complexation reactions from Table 24. Error bars that are not visible are smaller than the symbol size.

Figure 42 .

 42 Figure 42. Sorption isotherms of U(VI) onto illite in the presence of carbonates in equilibrium with atmospheric pCO2, 1% and 5% of CO2(g) in 0.1 NaNO3, pH fixed at 7.3 (• and •), and 6.6 (•), respectively. Error bars that are not visible are smaller than the symbol size.

Figure 43 .

 43 Figure 43. Model approach for the sorption isotherms of U(VI) onto illite in 0.1 M of NaNO3, S/L of 2 g/L at pH 7.3 in equilibrium with 1% of CO2(g) (left top and bottom) and pH 6.6 in equilibrium with 5% of CO2(g) (right top and bottom). Model II: only ≡S s OUO 2 (CO 3 ) 2 3-in considered; Model III: U(VI)-CO3 surface complexation reactions from Table24. Error bars that are not visible are smaller than the symbol size.

Figure 44 .

 44 Figure 44. Sorption isotherms of U(VI) on illite in the presence of 1% of CO2(g) under 0.1 M NaNO3 at pH 7.3 and synthetic COx porewater at pH 7.1. Error bars that are not visible are smaller than the symbol size.

  ) are shown in Figure 45. The model II can reproduce well the uptake of U(VI) onto illite in the presence of Ca-U(VI)-CO3 ternary complexes.

Figure 45 .

 45 Figure 45. Sorption isotherms of U(VI) on illite as a function of the U(VI) in COx porewater in equilibrium with 1% of CO2(g). Model II: only ≡ 𝑺 𝒔 𝑶𝑼𝑶 𝟐 (𝑪𝑶 𝟑 ) 𝟐 𝟑-in considered taking also in account the thermodynamic data for Ca-U(VI)-CO3 ternary complexes (see Chapter 2). Error bars that are not visible are smaller than the symbol size.

Figure 46

 46 Figure 46. a) Sorption isotherm of U(VI) onto (o) illite and (◊) Callovo-Oxfordian clay fraction (illite + I/S) in the absence of carbonates in both 2 g/L under 0.1 M NaNO3 at pH of 6.4 and 7.0 (respectively). b) Sorption isotherm of U(VI) onto (•) illite and (♦) Callovo-Oxfordian clay fraction (illite + I/S) in the presence of carbonates in equilibrium with 1% of CO2(g) in both 2 g/L under 0.1 M NaNO3 at pH of 7.3 and 7.2 (respectively). Model approach for illite (solid-line) and for COx clay fraction (dash-line). Error bars that are not visible are smaller than the symbol size.

Figure 47 .

 47 Figure 47. a) Sorption isotherms of U(VI) on 2 g/L of COx clay fraction (illite + I/S) in equibilibrium with 1% of CO2(g) in 0.1 M of NaNO3 (♦) and synthetic COx porewater (♦). b) Sorption isotherm of U(VI) on (•) 10 g/L of illite and (♦) on 2 g/L of COx clay fraction (illite + I/S) in equilibrium with 1% of CO2(g) in synthetic COx porewater conditions. Model approach for illite (solid-line) and for COx clay fraction (dash-line). Error bars that are not visible are smaller than the symbol size.

Figure 48 .

 48 Figure 48. Sorption isotherm of U(VI) onto (•) illite and (♦) COx clay fraction (illite + I/S) and (□) Callovo-Oxfordian clays sample in equilibrium with 1% of CO2(g) in synthetic COx porewater conditions. Error bars that are not visible are smaller than the symbol size.

Figure 49 .

 49 Figure 49. Model approach (solid line) for sorption isotherm of U(VI) on Callovo-Oxfordian claystone samples in equilibrium with pCO2 of 1% of CO2(g) in Callovo-Oxfordian synthetic porewater with pH 7.5. Error bars that are not visible are smaller than the symbol size.

Figure 50 .

 50 Figure 50. Experimental and model (black line) data from Hartmann et al. (2008) and the prediction model proposed in the present work (red lines) considering the new stability constant ≡S s OUO 2 (CO 3 ) 2 3-and its respective uncertainties.

Figure 51 .

 51 Figure 51. Comparison of sorption isotherms of U(VI) on the Opalinus clay from [1] Joseph et al. (2011) and Boda clay [2] Marques Fernandes et al. (2015) and [3] Amayri et al. (2016) with the COx clayey rock from the [p.w.] present work.

Figure 52 .

 52 Figure 52. Model prediction for the sorption isotherms of U(VI) on a) OPA clay from [1] Joseph et al. (2011); b) OPA clay from [2] Marques Fernandes et al. (2015); and c) Boda clay from [2] Marques Fernandes et al. (2015); [p.w.] COx clay from present work. Solid-lines: model approach proposed in present work; Dash-line: model uncertainties of the ≡S s OUO 2 (CO 3 ) 2 3-(log10K = 17.5 ± 0.5).

Figure 53 .

 53 Figure 53. Impact of the temperature on the distribution coefficient of U(VI) as a function of the concentration of U(VI) in equilibrium in COx claystone samples. (•) and (□) experimental data from the present work. red-line: model approach at 80ºC considering the model approach proposed in Chapter 3 and the thermodynamic parameters of U(VI) aqueous speciation (Chapter 2).

Figure 54 .

 54 Figure 54. Linear relationship between Log10Kºx of aqueous species and Log10 S Kºx of surface complexes for illite and montmorillonite, considering both hydrolysis complexes. Solid-line: correlation of illite surface complexation data; dashed-line: correlation of montmorillonite surface complexation data.

Figure 55 .

 55 Figure 55. Distribution coefficient for the sorption of U(VI) at 1•10 -7 M of U(VI) onto illite as a function of pH in 0.1 M of NaNO3 in the absence of carbonates. Error bars that are not visible are smaller than the symbol size.

Figure 56 .

 56 Figure 56. Temperature effect on the sorption isotherm of U(VI) onto illite in the absence of carbonates at pH of equilibrium of 6.4-6.5 in 0.1 M of NaNO3. Error bars that are not visible are smaller than the symbol size.

Figure 57 .

 57 Figure 57. Distribution of U(VI) aqueous species as a function of the pH at 20ºC (solid lines) and 80ºC (dashed lines) in the absence of the carbonates.

Figure 58 .

 58 Figure 58. Model approach for the distribution coefficient of U(VI) as a function of the pH at a) 20 ºC and b) 80 ºC in the absence of carbonates. Model II: base model without enthalpy data (see Chapter 3); Model IV: model considering the correlation factor 1 with aqueous and surface complexes for enthalpy values (Table33). Error bars that are not visible are smaller than the symbol size.

Figure 59 .

 59 Figure 59. Model approach on sorption of isotherms of U(VI) in the absence of carbonates at a) 20 ºC; b) 60 ºC; and c) 80 ºC. Model II: model approach without enthalpy values (see Chapter 3); Model IV: model approach considering the enthalpy values (see Table33). Error bars that are not visible are smaller than the symbol size.

Figure 60 .

 60 Figure 60. Distribution of U(VI) hydrolysis surface complexes of illite at 20ºC (solid-lines) and 80ºC (dashed-lines) as a function of the pH. Calculations taking the parameters given in Table33.

Figure 61 .

 61 Figure 61. Fitting model for sorption edge experiments (a) at 20 ºC and (b) at 80 ºC (with ±0.2 log units of uncertainty for the stability constants); for sorption isotherm experiments (c) at 20 ºC (with ±0.2 log units of uncertainty

Figure 62 .

 62 Figure 62. Relationship between Log10Kºx of aqueous species and Log10 S Kºx of surface complexes for illite determined at 20, 60 and 80ºC in comparison with the respective standard stability constants at 25ºC (solid-line) for illite reported by Bradbury and Baeyens (2009).

Figure 63 .

 63 Figure 63. Relation of constant of hydrolysis surface complexation (T) and standard constant of hydrolysis surface complexation (T0) as a function of the 1/T0-1/T for the surface species a) ≡S s OUO 2 + , b) ≡S s OUO 2 OH, c) ≡S s OUO 2 (OH) 2 -, and d) ≡S s OUO 2 (OH) 3 2-. Symbols: experimental data; Dashed-lines: regression line calculated for the enthalpy values.

Table 35 .Figure 64 .

 3564 Figure 64. Symbols: measured experimental data from sorption edge and isotherm of U(VI) at different temperatures. Solid-lines: calculated U(VI) concentration using enthalpy values calculated from Table 35. Dashed-lines

Figure 65 .

 65 Figure 65. Distribution coefficient of U(VI) as a function of the pH in 0.1 NaCl a) in the absence or presence of carbonates in equilibrium with pCO2 atmospheric at 80 ºC. Error bars that are not visible are smaller than the symbol size.

Figure 66 .

 66 Figure 66. Sorption isotherms of U(VI) on illite at different temperatures in equilibrium with a) Ar(g); b) pCO2 atmospheric; c) pCO2 of 1% of CO2(g); and d) pCO2 of 5% of CO2(g). Error bars that are not visible are smaller than the symbol size.

Figure 67 .

 67 Figure 67. Distribution of the speciation U(VI) in solution as a function of the pH at 20 (solid-lines) and 80 ºC (dashed-lined) in equilibrium with pCO2 atmospheric.

Figure 68 .

 68 Figure 68. Sorption edge of U(VI) in equilibrium with pCO2 atmospheric. Model V: enthalpy values considered only for hydrolysis surface complexation reactions (Table35).

Figure 69 .

 69 Figure 69. Model approach considering only the enthalpy of the hydrolysis of surface complexes at different temperatures in the presence of carbonates (see Table 35). Dashed-lines correspond to the uncertainty of the enthalpy values of the hydrolysis surface complexes.

Figure 70 .

 70 Figure 70. Sorption isotherms of U(VI) on illite and Callovo-Oxfordian clay fraction (illite + I/S) system at 80 ºC a) in the absence of carbonates and b) in the presence of carbonates in equilibrium with pCO2 of 5% of CO2(g).

Figure 71 .

 71 Figure 71. Model approach for describing the sorption isotherms of U(VI) on Callovo-Oxfordian clay fraction (illite + I/S) at 20 and 80 ºC a) in the absence and b) in the presence of carbonates in equilibrium with 1% and 5% of CO2(g), respectively. Taking the thermodynamic parameters for illite and montmorillonite (see Table 39) the solidlines correspond to the model prediction at 20ºC; and the dashed-lines correspond to the model prediction at 80 ºC.

Figure 72 .

 72 Figure 72. Sorption isotherms of the U(VI) in synthetic porewater performed at 20 and 80 ºC for a) U(VI)/illite system; b) U(VI)/ Callovo-Oxfordian clay fraction system; and c) U(VI)/ Callovo-Oxfordian claystone system. Model approach considering thermodynamic data for surface reactions from Table 39 and using the thermodynamic data for Ca-U(VI)-CO3 aqueous complexes (see Table 14). Solid-line: model at 20 ºC; dashed-line: model at 80ºC.

Figure 73 .

 73 Figure 73. Impact of the temperature on the distribution coefficient of U(VI) as a function of the concentration of U(VI) in equilibrium in Callovo-Oxfordian claystone system. (•) and (□) experimental data from the present work.

  corresponds to the standard deviation of the ycalc, m to the number of measurements unknown, and n the number of calibration points.The limite of detection was calculated as: = 𝑘 • 𝑠𝑥, where k = 2 for 95% of significance.

Figure A. 1 .

 1 Figure A. 1. Schematic representation of the calibration plot of ICP-MS measurements. Black-cross correspond to the measurement of each standard solution; solid-line corresponds to the regression line (𝒚 𝒄𝒂𝒍𝒄 = 𝒃𝒙 + 𝒄); dash-lines corresponds to the inferior and superior error of ycalc.

  The Callovo-Oxfordian claystone was purified, and the clay fraction was separated. The separation and purification of the clay fraction was performed and provided by BRGM. The clay fraction samples were storage in open atmosphere.The procedure is divided in 4 stages: Stage 1: separation of the < 2 µm fraction; Stage 2: acidification of the carbonates; Stage 3: Oxidation of iron and manganese oxides; and Stage 4: reduction of organic matter. The procedure of separation and purification is described below:Stage 1: < 2 µm Clay Fraction Separation 1. 20 g/L clay dispersed in Milli-Q Water.2. The fine fraction (<2 µm) separated by sedimentation or centrifugation 3. The depth of fine fraction (<2 µm) is calculated by using Stoke's Law. 4. Fine fraction is taken into vacuum flask 5. Centrifuge suspension at high speed to get 0.1-2 µm clay fraction Stage 2: Carbonate Removal with Acetic Acid (treated with 0.1 M acetic acid in 0.5 M NaCl (pH ~5)) 1. Prepare 2 L 0.2 M Acetic acid stock solution at pH 5 2. 20 g clay fraction (dry equivalent) dispersed in 0.5 L Milli-Q water. 3. 0.5 L clay suspension mixed with 0.5 L 0.2 M Acetic acid (pH ~5) in a pot 4. Add 29.22 g NaCl into pot and close the pot, shake by hand few seconds and open the pot in order to release excess CO2 pressure under a vapor extractor 5. Close the pot and shake it on horizontal shaker for 1 h 6. High speed centrifugation to settle particles above 0.1 μm 7. Waste the supernatant (!! Not in the sink; organic acid) 8. Repeat acid attack one more time 9. After last centrifugation, disperse the clay with 0.5 M NaCl solution, shake and centrifuged 10. Repeat washing 2 times Stage 3: Oxidation to remove Fe and Mn oxides with DCB (dithionite-citrate-bicarbonate) (treated with 0.2 M Citrate, NaCl 0.5 M, NaHCO3 0.1 M (pH ~6)) 1. Prepare 2 L stock solution having 0.4 M Citrate, 1 M NaCl, 0.2 NaHCO3 M (pH ~6) First solve citrate and buffered with NaOH, then add NaCl and NaHCO3 then final pH become ~6 2. In order to get fine dispersed clay suspension, 20 g (dry equivalent) clay fraction obtained from stage 2 dispersed in 0.5 L water. 3. 0.5 L clay suspension mixed with 0.5 L buffered Citrate stock solution in 1 L pot 4. Add one spatula (teaspoon) of sodium dithionite into pot and closed it. 5. Shake and put the pots at 40°C for 5 minutes and then open the pot under the vapor extractor 6. Close the pot and shake it on horizontal shaking table for 1 h 7. Centrifuge the suspension 8. Repeat DCB attack 1 more time 9. After last centrifugation, wash the clay by dispersing with 0.5 M NaCl solution, shake and centrifuge 10. Repeat washing 2 times Stage 4: Reduction to remove organic matter (treated with 3% H2O2, 0.5 M NaCl at 60 °C ) 1. Prepare 2 L stock solution H2O2 6% + NaCl 1 M 2. 20 g (dry equivalent) clay dispersed clay in 0.5 L water 3. Rinse a big glass beaker with the H2O2 solution 4. Add 0.5 L clay dispersion and 0.5 L H2O2 6% + NaCl 1 M Solution 5. Pour the clay dispersion in the beaker and heat the beaker at 55°C-60°C (STAY IN FRONT OF THE BEAKER AND WATCH FOR THE TEMPERATURE), 1 h 6. Wait for cooling, put in centrifuge pot and centrifuge the clay (use old but not cracked pots) 7. Wash clay by dispersing with 0.5 M NaCl solution, shake and centrifuged 8. Repeat washing 2 more times

  Figure C. 2. Temperature measurements as a function of time at open lid and closed lid with small aperture.

Figure C. 3 .

 3 Figure C. 3. Temperature measurements as a function of time in presence or absence of the isolation Teflon layer.

  Figure C. 4. Experimental test of the oxidation of Fe(II) to Fe(III) as gas sealing control.

Figure C. 5 .

 5 Figure C. 5. pH measured in atmospheric conditions of the COx porewater as a function of time.

Figure C. 7 .

 7 Figure C. 7. Comparison results of the number of washes with ultra-pure water between 3-4 days or 1 days.

Figure C. 8 .

 8 Figure C. 8. Total organic carbon as function of the munber of washing and time of contact. a) washing with ultrapure water and rinsed with ultra-pure water between each wash; b) first rinsed with 5% of HNO3 at 100 ºC before washed with ultra-pure water and rinsed with ultra-pure water between washes.

Figure

  Figure D. 2. (left) distribution of U(VI) speciation and (right) saturation index of Calcite as a function of the calcium concentration considering the experimental conditions of Bernhard et al. (2001) using Thermochimie database Davies (2014). Solid-line: U(VI) = 1•10 -4 M; dash-line: U(VI) = 5•10 -5 M; dot-line: U(VI) = 2•10 -5 M.

Figure D. 3 .

 3 Figure D. 3. Distribution of U(VI) speciation and the saturation index of calcite as a function of the calcium concentration considering the experimental conditions of Lee and Yun (2013) using Thermochimie database Davies (2014).

Figure D. 4 .

 4 Figure D. 4. (left) distribution of U(VI) speciation and (right) mineral saturation index as a function of the calcium concentration considering the experimental conditions of Endrizzi and Rao (2006) using Thermochimie database Davies (2014).

Figure

  Figure D. 5. Distribution of U(VI) speciation at 20ºC and 1% of CO2(g) taking the log10β113 = 27.4 ± 0.1 and log10β213 = 29.8 ± 1.1 (calculated from PEST-Phreeqc).

Figure

  Figure D. 6. Distribution of U(VI) speciation at 40ºC and 1% of CO2(g) taking the log10β113 = 27.2 ± 0.1 and log10β213 = 29.9 ± 0.7 (calculated from PEST-Phreeqc).

Figure

  Figure D. 7. Distribution of U(VI) speciation at 60ºC and 3% of CO2(g) taking the log10β113 = 26.7 ± 0.03 and log10β213 = 29.6 ± 0.1 (calculated from PEST-Phreeqc).

Figure

  Figure D. 8. Distribution of U(VI) speciation at 80ºC and 5% of CO2(g) taking the log10β113 = 26.5 ± 0.1 and log10β213 = 29.7 ± 0.1 (calculated from PEST-Phreeqc).

  the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: TOT("U") =0 then Kd = 1e-10 else Kd = (SYS("U") -TOT("U"))/TOT("U")/wt 10 U_ads = SYS("U") 20 U_sol = TOT("U") 30 wt = 3.94e-2 40 KD = (U_ads -U_sol)/U_sol/wt #50 ratio = LOG10(KD) 100 SAVE log10(KD) (SYS("U") -TOT("U"))/SYS("

  the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium Titre : Impact de l'élévation de la température jusqu'à 80 °C sur le comportement des radionucléides dans le Callovo-Oxfordien : application à l'uranium Mots clés : Ca-U(VI)-CO3 ; sorption ; température ; illite ; les argiles du Callovo-Oxfordien ; déchets radioactifs Résumé : Ce travail vise à comprendre et quantifier le comportement de U(VI) dans les argilites du Callovo-Oxfordien (COx), prévues pour accueillir les déchets nucléaires de haute activité en France. L'effet de la température sur ce comportement est particulièrement étudié. La première partie du travail s'est concentrée sur les propriétés thermodynamiques des complexes ternaires Ca-U(VI)-CO3 qui contrôlent la spéciation de U(VI) en solution. Ces dernières ont été mesurées par une méthode de compétition en présence d'une résine sous pCO2 et pH contrôlés. Les résultats indiquent que la température ne favorise pas la formation de CaUO2(CO3)3 2-(log10 βº113 = 27,3 ± 0,3 ; ΔrHº = -27,4 ± 8 kJ/mol) et n'affecte pas la formation de Ca2UO2(CO3)3(aq) (log10 βº213 = 29,7 ± 0,3 ; ΔrHº = 0 ± 2 kJ/mol). Une approche « bottom-up » avec le modèle « 2SPNE SC / CE » publié dans la littérature a été suivie pour décrire les phénomènes de sorption, en considérant que la fraction argileuse du COx (Illite et I/S) gouverne la sorption de U(VI). Ce modèle a été appliqué avec succès pour reproduire une multitude de données expérimentales obtenues avec l'illite, la fraction argileuse du COx et les argilites du COx en fonction de paramètres clés (pH, pCO2, [U(VI)], [Ca]) à 20 °C. Le phénomène de rétention dans les conditions in-situ est principalement régi par la sorption des complexes U(VI)carbonate et une nouvelle constante d´équilibre de réaction de complexation de surface est proposée pour l'illite. Une augmentation de la température à 80 °C conduit à une augmentation de la rétention de U(VI) sur le COx. Cette augmentation est accompagnée d'un changement de pCO2 et de pH. Le modèle de rétention testé à 20 °C combiné avec les paramètres thermodynamiques décrivant le comportement de U(VI) en solution expliquent cette augmentation sans pour autant obtenir un accord satisfaisant avec l'expérience. Le modèle est amélioré en intégrant des valeurs de ΔrHº obtenues pour les réactions de complexation de surface à partir du système U(VI)/illite. Title: Impact of the increase in temperature up to 80 ºC on the behaviour of radionuclides in the Callovo-Oxfordian formation: application to uranium Keywords : Ca-U(VI)-CO3; sorption; temperature; illite; Callovo-Oxfordian claystone; radioactive waste Abstract : The aim of this study was to understand and quantify the behaviour of U(VI) on the Callovo-Oxfordian (COx) clay which is envisioned to host high-level radioactive waste in France. The temperature effect up to 80°C on this behaviour was particularly studied. The first part of the work focussed on the thermodynamic properties of the calcium uranyl carbonate aqueous complexes which govern U(VI) speciation in solution. They were measured indirectly by sorption-based methodologies under controlled pCO2 and pH. The results indicate that the temperature does not favour the formation of CaUO2(CO3)3 2-(log10 βº113 = 27.3 ± 0.3; ΔrHº = -27.4 ± 8 kJ/mol) and does not affect the formation of Ca2UO2(CO3)3(aq) (log10 βº213= 29.7 ± 0.3; ΔrHº = 0 ± 2 kJ/mol).

Table 2 .

 2 . ................................................................................................ 34 Chemical composition of the Callovo-Oxfordian porewater calculated by geochemical modelling according to the temperature increasing using the reduced mineralogy. ............................................................................. 35

Table 3 .

 3 Chemical composition of the Callovo-Oxfordian porewater calculated by geochemical modelling according to the temperature increasing using the completed mineralogy. ......................................................................... 35

Table 4 .

 4 Summary of the literature analysing the stability constants of Ca-U(VI)-CO3 complexes. Data in grey have been superseded by later references of the same authors (2001). ....................................................................... 41

Table 5

 5

. Summary of the literature on the stability constants for Mg-U(VI)-CO3 complexes. Data in grey have been superseded by later references of the same authors. ........................................................................................... 42

Table 6. Stability constants for Ca-U(VI)-CO3 complexes at I = 0.1M and different temperatures suggested by Jo and Yun (2017). .................................................................................................................................................. 43 Table 7. Summary of the batch experiments performed in the present work. .......................................................... 47 Table 8. Description of the tests performed in different temperatures and in closed system experiments. ............. 51 Table 9. Theoretical and measured parameters from the evaluation of CO2(g) in equilibrium at different temperatures. Theor.: theoretical value in agreement with the initial conditions of the experiments; Exp.: pH and alkanity measured at the end of the experiment, and pCO2 calculated according pH and alkanity measurements. ............................................................................................................................................................................ 51 Table 10. Estimated values for the Log10KX at different temperatures and pCO2 (atm). ......................................... 57 Table 11. Theoretical and measured pH, alkalinity and pCO2 of the batch experiments at different temperatures. Theor.: theoretical values according to the initial conditions of the experiments; Measured pH and alklinity in equilibrium; and calculated pCO2 in agreement with pH and alkalinity measured. ........................................... 58 Table 12. Results of the fitting for the stability constants of Ca-U(VI)-CO3 complexes using PEST. .................... 60 Table 13. Stability constants of Ca-U(VI)-CO3 complexes calculated by 2D mapping method. ............................. 62 Table 14. Thermodynamic data of Ca-U(VI)-CO3 ternary complexes determined in this work. ............................. 63 Table 15. Calculated UO2•2H2O(am) solubility and aqueous U(VI) speciation under Callovo-Oxfordian conditions (Tournassat et al., 2008, Table 3) ....................................................................................................................... 67

Table 16 .

 16 Comparison of the thermodynamic parameters of the Ca-U(VI)-CO3 ternary complexes reactions. p.w. present work; (1) Endrizzi and Rao (2017); (2) Jo and Yun (2017); n.d. not defined. ....................................... 67

Table 17 .

 17 Classification of the planar hydrated phyllosilicates and some examples for each clay mineral species (Theng, 2012). (the higlighted clay minerals are the ones relevant for the present study) ................................. 76

Table 18

 18

. Sorption parameters for U(VI) onto clay minerals of Callovo-Oxfordian formation

Table 20 .

 20 Distribution of surface complexation reactions of U(VI) on illite and montmorillonite in Callovo-Oxfordian formation system (pH = 7.2 and pCO2 = 10 -2.1 atm) .......................................................................... 84

Table 21 .

 21 Description of the experimental cases performed on different phases of Callovo-Oxfordian formation at 20 ºC. .................................................................................................................................................................. 85

Table 22 .

 22 Summary of the experimental conditions for sorption of uranium in different components of Callovo-Oxfordian formation. .......................................................................................................................................... 86

Table 23 .

 23 Alkalinity and pH measurements before and after the equilibrium with illite. pCO2 calculated from the equilibrium pH/alkalinity. ................................................................................................................................... 90

Table 24 .

 24 Uranyl carbonate surface complexation reactions proposed for illite in Model III. ................................. 92 Table 25. pH and alkalinity measurements after equilibrium with pCO2 atmospheric, 1% and 5% CO2(g) ........... 94

Table 26 .

 26 Surface species of U(VI) distribution on illite and montmorillonite for Callovo-Oxfordian formation experimental conditions. ................................................................................................................................... 100

Table 27 .

 27 Comparison of the main chemical parameters for OPA clay, Boda clay and COx clay. [1] Joseph et al.

(2011); [2] Marques Fernandes et al. (2015); [p.w.] present work. (*) measurements of alkalinity (meq/L). . 102

Table 28. Compilation of the surface reactions for both illite and montmorillonite taken in for the model prediction. (n.a. not applied) ............................................................................................................................................... 104 Table 29. Description of the U(VI) sorption experiments carried out on illite, Callovo-Oxfordian clay fraction and clays samples as a function of temperature. ...................................................................................................... 113 Table 30. Experimental conditions of the sorption experimental cases for the different phases of Callovo-Oxfordian formation at different temperatures. (pw = porewater). .................................................................................... 113 Table 31. Aqueous speciation of U(VI)

Table 37 .

 37 pH of equilibrium and alkalinity measurements in the presence of carbonates at 60 and 80ºC. ............ 126 Table 38. pH and alkalinity measurements in equilibrium for the sorption isotherms on clay fraction at 80 ºC in different atmospheres. ....................................................................................................................................... 129

. (n.a.: not applied). .............................................. 114

Table 32. pH and alkalinity measured in equilibrium for the sorption isotherms on illite in inert atmosphere at different temperatures. ...................................................................................................................................... 117 Table 33. Surface complexation constants on strong sites together with the corresponding hydrolysis constants and enthalpy values for illite considered in Model IV. ............................................................................................ 118 Table 34. Log10 OH Kxº fitted based on sorption edges and isotherms experiments performed at different temperatures for illite. ............................................................................................................................................................ 121 Table 35. Enthalpy values calculated from fitting of the stability constants of hydrolysis surface site species at different temperatures. The uncertainty was determined considering only the propagation error of the standard deviation of the stability constants. ................................................................................................................... 123 Table 36. Alkalinity and pH measurements before and after the equilibrium with illite at 80 ºC. pCO2 calculated from the equilibrium pH/alkalinity. .................................................................................................................. 125

Table 39 .

 39 Enthalpy values for the hydrolysis surface complexation reactions for illite and montmorillonite considered in the present study. ........................................................................................................................ 129

Table 40 .

 40 Experimental pH and alkalinity measurements at 20 and 80 ºC for illite, Callovo-Oxfordian clay fraction and claystone samples in synthetic Callovo-Oxfordian porewater in equilibrium with 1% and 5% of CO2(g),

respectively. ...................................................................................................................................................... 130 Table 41. Aqueous and surface speciation at 20 and 80ºC for the Callovo-Oxfordian formation system in equilibrium with synthetic porewater for [U(VI)] = 10 -7 M. ................................................................................................. 132 List Of Tables xix Flávia Maia | IMT Atlantique Nantes

Table 42 .

 42 Proposed thermodynamic parameters for the surface complexation reactions considered in the model approach. ........................................................................................................................................................... 133 Table A. 1. Concentration and dilution calculations of the standard solutions for determination of the calibration curve of ICP-MS analysis ................................................................................................................................. 142 Table A. 2. Theoretical chemical composition of the Callovo-Oxfordian porewater at 20 ºC and 80ºC. ............. 146 Table A. 3. Salt content in g/L used in preparation of synthetic COx porewater at 20 and 80ºC ......................... 147 Table A. 4. Mineral composition of Callovo-Oxfordian claystone ....................................................................... 147 Table A. 5. Experimental conditions for the batch experiments. Temp.= temperature; COx = Callovo-Oxfordian; I = Ionic strength. ................................................................................................................................................ 150 Table B. 1. Amount of exchanger-NO3 (mol/L) for the different solid-liquid ratio (𝝌) considered in the closed system experiments. .

	Table C. 12. Experimental results of the ionic exchange method at 20ºC in closed system at 1% of CO2(g) in the
	presence of calcium, 0.1 M of NaNO3 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M. (blank
	samples in the absence of calcium (0) or presence of calcium (19) for Ba: without U(VI); and Bb: without resin
	with U) .Table C. 13. Experimental results of the ionic exchange method at 40ºC in closed system at 1% of CO2(g) in the
	absence of calcium and varying the solid-liquid ration, in 0.1M of NaNO3 adding initially [U(VI)] = 5•10 -7 M.

......................................................................................................................................... 153

Table C. 1. Experimental test of the control of the temperature using an isolation-aluminium box without external layer of Teflon. ................................................................................................................................................. 156 Table C. 2. Experimental test of the control of the temperature using an isolation-aluminium box with external layer of Teflon. .......................................................................................................................................................... 157 Table C. 3. Experimental test for sealing control. ................................................................................................. 159 Table C. 4. Experimental results of the tightness of the tubes without sealing tape. ............................................ 160 Table C. 5. Uncertainty of the signal of the TOC Analyser determined using standard solution of IC, TC and ultrapure water. ........................................................................................................................................................ 162 Table C. 6. TOC measurements of the experimental tests. ................................................................................... 163 Table C. 7. Chemical composition of the synthetic Callovo-Oxfordian porewater. Uncertainty of k=2 .............. 164 Table C. 8. Measurements of the pH for the sorption edge experiments in atmospheric pCO2. (1) used only in the experiments at 80ºC; (2) this solution was not taken for the experiments performed at 20ºC. .......................... 164 Table C. 9. Alkalinity measurements using H2SO4 (0.16 N) in the blank samples of the batch experiments using anionic-exchange resin. *measured .................................................................................................................. 164 Table C. 10. Experimental results of the ionic exchange method at 20ºC in atmospheric conditions in the presence of calcium, 0.1 M of NaClO4 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M............................ 165 Table C. 11. Experimental results of the ionic exchange method at 20ºC in closed system at 1% of CO2(g) in the absence of calcium and varying the solid-liquid ration, in 0.1M of NaNO3 adding initially [U(VI)] = 5•10 -7 M. (Ba: blank sample without U(VI); Bb: blank sample without resin with U) .................................................... 165 ............................................................................................................................................................. 166 (Ba: blank sample without U(VI); Bb: blank sample without resin with U) .................................................... 166 Table C. 14. Experimental results of the ionic exchange method at 40ºC in closed system at 1% of CO2(g) in the presence of calcium, 0.1 M of NaNO3 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M. . (blank samples in the absence of calcium (0) or presence of calcium (19) for Ba: without U(VI); and Bb: without resin with U) .............................................................................................................................................................. 166

Table C .

 C 15. Experimental results of the ionic exchange method at 60ºC in closed system at 3% of CO2(g) in the absence of calcium and varying the solid-liquid ration, in 0.1M of NaNO3 adding initially [U(VI)] = 5•10 -7 M. (Ba: blank sample without U(VI); Bb: blank sample without resin with U) .................................................... 167 Table C. 16. Experimental results of the ionic exchange method at 60 ºC in closed system at 3% of CO2(g) in the presence of calcium, 0.1 M of NaNO3 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M. . (blank samples in the absence of calcium (0) or presence of calcium (19) for Ba: without U(VI); and Bb: without resin with U) .............................................................................................................................................................. 167 Table C. 17. Experimental results of the ionic exchange method at 80 ºC in closed system at 5% of CO2(g) in the absence of calcium and varying the solid-liquid ration, in 0.1M of NaNO3 adding initially [U(VI)] = 5•10 -7 M. (Ba: blank sample without U(VI); Bb: blank sample without resin with U) .................................................... 168 Table C. 18. Experimental results of the ionic exchange method at 80ºC in closed system at 5% of CO2(g) in the presence of calcium, 0.1 M of NaNO3 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M. . (blank samples in the absence of calcium (0) or presence of calcium (19) for Ba: without U(VI); and Bb: without resin with U). ............................................................................................................................................................. 168

Table C

 C

	Table C. 31. Experimental results of the sorption isotherms of U(VI) onto illite at 60ºC in 1% of CO2(g) in 0.1 M
	of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 178
	Table C. 32. Experimental results of the sorption isotherms of U(VI) onto illite at 60ºC in 5% of CO2(g) in 0.1 M
	of NaNO3 at fixed solid-liquid ratio.

. 19. Kinetic results of sorption of U(VI) onto Callovo-Oxfordian claystone at 20 ºC in 1% CO2(g), using synthetic Callovo-Oxfordian porewater and initial U(VI) = 6.8(±0.2)•10 -7 M. ................................................. 168

Table C. 20. Kinetic results of sorption of U(VI) onto Callovo-Oxfordian claystone at 80 ºC in 5% CO2(g), using synthetic Callovo-Oxfordian porewater and initial U(VI) = 6.8(±0.2)•10 -7 M. ................................................. 169 Table C. 21. Kinetic results of sorption of U(VI) onto Illite at 20 ºC in 1% CO2(g), using synthetic Callovo-Oxfordian porewater and initial U(VI) = 9.6(±0.8)•10 -8 M. ............................................................................... 169 Table C. 22. Kinetic results of sorption of U(VI) onto illite at 80 ºC in 5% CO2(g), using synthetic Callovo-Oxfordian porewater and initial U(VI) = 6.8(±0.2)•10 -7 M. ............................................................................... 169 Table C. 23. Alkalinity measurements in sorption isotherms experiments of U(VI)/illite system (2 g/L) at different temperatures. *synthetic Callovo-Oxfordian porewater. .................................................................................. 170 Table C. 24. Experimental results of the sorption isotherms of U(VI) onto illite at 20ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................................. 171 Table C. 25. Experimental results of the sorption isotherms of U(VI) onto illite at 20ºC in atmospheric conditions in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) .... 172 Table C. 26. Experimental results of the sorption isotherms of U(VI) onto illite at 20ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 173 Table C. 27. Experimental results of the sorption isotherms of U(VI) onto illite at 20ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 174 Table C. 28. Experimental results of the sorption isotherms of U(VI) onto illite at 20ºC in 1% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ............................................................................................................................................................ 175 Table C. 29. Experimental results of the sorption isotherms of U(VI) onto illite at 60ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................................. 176 Table C. 30. Experimental results of the sorption isotherms of U(VI) onto illite at 60ºC in atmospheric pCO2 in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)............... 177 (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 179 Table C. 33. Experimental results of the sorption isotherms of U(VI) onto illite at 80ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................................. 180 Table C. 34. Experimental results of the sorption isotherms of U(VI) onto illite at 80ºC in atmospheric pCO2 in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)............... 181 List Of Tables xxi Flávia Maia | IMT Atlantique Nantes

Table C .

 C 35. Experimental results of the sorption isotherms of U(VI) onto illite at 80ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 182 Table C. 36. Experimental results of the sorption isotherms of U(VI) onto illite at 80ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................ 183 Table C. 37. Experimental results of the sorption isotherms of U(VI) onto illite at 80ºC in 5% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ............................................................................................................................................................ 184 Table C. 38. pH results of each tube measured before the batch equilibrium at 20 ºC in atmospheric pCO2. ...... 185 Table C. 39. pH results of each tube measured before the batch equilibrium at 80 ºC in atmospheric pCO2. ...... 185 Table C. 40. Experimental results of the sorption edges of U(VI) onto illite at 20ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio of 2 g/L with [U(VI)]initial = 1.02(±0.06)•10 -7 M. (BF: blank sample filtrated; BNF: blank sample no filtrated) ........................................................................................................................................... 186

	Table C. 41. Experimental results of the sorption edges of U(VI) onto illite at 20ºC in atmospheric pCO2 in 0.1 M
	of NaCl at fixed solid-liquid ratio of 0.5-1 g/L with [U(VI)]initial = 1.02(±0.06)•10 -7 M. (BF: blank sample filtrated;

BNF: blank sample no filtrated) ....................................................................................................................... 186

Table C. 42. Experimental results of the sorption edges of U(VI) onto illite at 80 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio of 2 g/L with [U(VI)]initial = 1.02(±0.06)•10 -7 M. (BF: blank sample filtrated; BNF: blank sample no filtrated) ........................................................................................................................................... 187 Table C. 43. Experimental results of the sorption edges of U(VI) onto illite at 80ºC in atmospheric pCO2 in 0.1 M of NaCl at fixed solid-liquid ratio of 0.5-1 g/L with [U(VI)]initial = 3.07(±0.01)•10 -8 M. (BF: blank sample filtrated; BNF: blank sample no filtrated) ....................................................................................................................... 187 Table C. 44. Alkalinity measurements in sorption isotherms experiments of U(VI)/clay fraction system at different temperatures. *synthetic Callovo-Oxfordian porewater. .................................................................................. 188 Table C. 45. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. .................................................................................... 188 Table C. 46. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20 ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ....................................................................................................................................................... 189 Table C. 47. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20 ºC in 1% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ....................................................................................................................... 190 Table C. 48. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio of 2 g/L. (BF: blank sample filtrated; BNF: blank sample no filtrated) ....................................................................................................................................................... 191 Table C. 49. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ....................................................................................................................................................... 192 Table C. 50. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in 5% of CO2(g) in Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated) ................................................................................................................................. 193 Table C. 51. Alkalinity measurements in sorption isotherms experiments of U(VI)/claystone system at 20 and 80ºC in synthetic Callovo-Oxfordian porewater. ....................................................................................................... 194 Table C. 52. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian claystone at 20 ºC in 1% of CO2(g) in Cox porewater at fixed solid-liquid ratio. .............................................................................. 194 Table C. 53. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian claystone samples at

Table D .

 D 1. Summary of the main parameters used for determination of stability constants of Ca-U(VI)-CO3 complexes using different experimental methods (n.d.: not determined). ........................................................ 198 Table D. 2. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 20 ºC and 1% of CO2(g) ...................................................................................................................... 200 Table D. 3. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 40 ºC and 1% of CO2(g) ...................................................................................................................... 200 Table D. 4. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 60 ºC and 3% of CO2(g) ...................................................................................................................... 200 Table D. 5. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 80 ºC and 5% of CO2(g) ...................................................................................................................... 201 Table D. 6. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 20ºC and 1% of CO2(g) ..................................................................................................................................... 201 Table D. 7. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L,

40ºC and 1% of CO2(g) ..................................................................................................................................... 202

Table D. 8. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 60ºC and 3% of CO2(g) ..................................................................................................................................... 202 Table D. 9. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 80ºC and 5% of CO2(g) ..................................................................................................................................... 203 Table E. 1. Thermodynamic data for the aqueous species considered for the calculation of stability constants of Ca-U(VI)-CO3 complexes. Note that log10β values and enthalpies of reaction are in agreement with the chemical reactions that are included in ThermoChimie database v.9 .............................................................................. 205 Table E. 2. Thermodynamic data for the solid phases considered for the calculation of stability constants of Ca-U(VI)-CO3 complexes. Note that log10K and enthalpies of reaction are in agreement with the chemical reactions that are included in ThermoChimie database v.9. ............................................................................................. 206

Modèle de sorption U(VI)/argilites du Callovo-Oxfordien à 20 °C

	CaUO2(CO3)3 2-et Ca2UO2(CO3)3(aq) sont respectivement égales à -27,4 ± 8,2 et à 0 ± 1,8 kJ/mol. La
	valeur de ΔrHº du complexe anionique est négative, comme cela est proposé par Endrizzi et Rao
	(2014), même si celle-ci est deux fois plus faible dans le cas présent. Cependant, pour le complexe
	neutre, la valeur de ΔrHº est en désaccord avec celle proposée par Endrizzi et Rao (2014) (-47 kJ /
	mol). Les autres paramètres thermodynamiques (entropie et énergie libre de Gibbs) sont en bon ac-
	cord avec ceux publiés par Endrizzi et Rao (2014).
	La formation des argilites du Callovo-Oxfordien est constituée principalement de quartz, de minéraux
	carbonatés (comme la calcite, la dolomite), de pyrite et d'une fraction argileuse qui contient en ma-
	jorité de l'illite et des argiles interstratifiées I/S (illite/smectite). Pour simplifier l'étude, dans une
	première approche, les hypothèses suivantes sont faites : (i) la fraction argileuse est considérée
	comme étant à l'origine de la rétention de l'uranium ; (ii) la fraction I/S est décrite comme un mélange
	d'illite et de smectite et enfin (iii) la réactivité de la phase argileuse est décrite par deux pôles qui sont
	l'illite et la montmorillonite (comme minéral représentatif des smectites)
	2-(log10 βº113
	= 27,3 ± 0.2) et Ca2UO2(CO3)3(aq) (log10 βº213 = 29.7 ± 0.3) ont été calculées en suivant deux ap-
	proches d'ajustement différentes. Les valeurs obtenues sont cohérentes avec celles publiées, ce qui
	valide la méthode et le modèle conceptuel proposé. De plus, les données d'enthalpie ont été détermi-
	nées pour la formation des complexes ternaires Ca-U(VI)-CO3. Les enthalpies de formation de

Modèle de sorption U/argilites du Callovo-Oxfordien en température (20 à 80 °C)

	alors que la spéciation à la surface est dominée par le complexe ≡SsOUO2(CO3)3 2-. Au-delà de ce
	schéma simplifié, il est important de noter que le modèle prend en compte un nombre important
	d'équilibres, tels que les réactions d'échange cationique ou celles qui décrivent les propriétés acido-
	basiques des sites de surface. Il est donc difficile d'évaluer l'incertitude du modèle par rapport à ces
	capacités prédictives, bien que des études préliminaires de sensibilité indiquent que l'incertitude sur
	le complexe carbonaté de surface influence fortement les résultats de calcul dans les conditions du
	COx.
	Le comportement de la sorption de U(VI) en température a été étudié en partant du système le plus
	simple (illite) vers le système le plus complexe (assemblage minéralogique argileux de la formation
	du Callovo-Oxfordien) en suivant l'approche mise en place pour l'étude à 20 °C. Les paramètres clés
	étudiés sont le pH, la pression partielle de CO2(g) et la composition de la solution.
	L'une des principales conclusions de cette étude est que l'effet de la température sur la sorption de
	U(VI) sur l'illite est négligeable en présence de NaNO3 ou NaCl ; les valeurs de Kd obtenues dans
	des conditions expérimentales similaires (pH, force ionique…) ne sont pas influencées dans la gamme
	de température étudiée (20-80 °C). Ceci est en contradiction avec l'analyse quantitative des données.
	En effet, la formation des complexes aqueux (et notamment les complexes d'hydrolyse) est favorisée
	lorsque la température augmente (base de données ThermoChimie, Giffaut et al., 2014), ce qui devrait
	alors conduire à une diminution de la rétention. Pour permettre une interprétation quantitative des
	données, il est donc nécessaire de considérer également un effet endothermique pour décrire les phé-
	nomènes de sorption afin de contrebalancer l'effet observé en phase aqueuse. Les données ne peuvent
	pas être décrites en appliquant directement les valeurs d'enthalpie connues des espèces aqueuses aux
	espèces de surface correspondantes, car le modèle global surestime systématiquement les données
	expérimentales. Ainsi, les valeurs d'enthalpie pour les complexes de surface ont dû être ajustées à
	partir des données expérimentales en suivant une approche progressive en considérant un effet négli-
	geable de la température pour le mécanisme d'échange ionique et en prenant en compte deux con-
	traintes pour les complexes de surfaces : l'équation de Van't Hoff doit être respectée et une corréla-
	tion linéaire doit s'appliquer entre les valeurs pour la spéciation aqueuse et pour la spéciation de
	surface. Un modèle de complexation de surface a ainsi pu être construit pour permettre d'expliquer
	de manière satisfaisante les données expérimentales. Dans le modèle, seules les valeurs d'enthalpie
	pour les espèces hydrolysées de surface ont dues être ajustées. Avec l'augmentation de la température,
	les espèces carbonatées de surface ne sont plus majoritaires pour les valeurs de pH inférieures à 7.
	Le modèle développé pour l'illite peut s'appliquer de manière satisfaisante aux échantillons naturels
	(fraction argileuse et échantillon brut des argilites du Callovo-Oxfordien) en absence ou en présence
	des complexes ternaires Ca-U(VI)-CO3.
	Le modèle proposé, avec les nouveaux paramètres obtenus, peut donc être utilisé pour évaluer le
	comportement de U(VI) dans la formation argileuse du Callovo-Oxfordien pour une gamme de tem-
	pérature variant entre 20 et 80 °C. Si l'uranium reste majoritairement sous le degré d'oxydation (VI)
	dans la solution porale du fait de la formation des complexes ternaires, la présence éventuelle de
	complexes de surface faisant intervenir l'uranium au degré d'oxydation U(IV) demeure une possibi-et de Boda (Marques Fernandes et lité. Cela peut être étudié si l'on est capable de contrôler expérimentalement le potentiel redox du al., 2015). Les résultats obtenus indiquent que la fraction argileuse gouverne la rétention de U(VI). De manière simplifiée, la spéciation en solution est contrôlée par les complexes ternaires de U(VI) système dans les conditions physico-chimiques représentatives des argilites à l'état naturel.

Table 1 . Average and extreme values of the mineralogical composition (wt%) of the host layer in borehole EST205 Meure/Haute-Marne site (Gaucher et al., 2004).

 1

	Minerals	wt%	max	min
	Illite group	12.9	21	2
	Interstratified illite-smectite (I/S)	11.6	33	0
	Kaolinite	1.7	5	0
	Chlorite	1.9	2.7	0
	Quartz (silicates)	9.5	43	0.9
	Calcite (carbonates)	15.3	75	2
	Pyrite	0.7	0.9	0.2
	Siderite + ferriferous minerals	1.5	2.7	0.5
	Titanium minerals	0.2	0.4	0
	Phosphate minerals	0.2	0.6	0.2
	1.1

.1.1. Chemical composition of the Callovo-Oxfordian porewater

Table 2 . Chemical composition of the Callovo-Oxfordian porewater calculated by geochemical modelling accord- ing to the temperature increasing using the reduced mineralogy.

 2

	Temp	pH	pCO2	Cl -	SO4 2+	TIC	Na +	K +	Ca 2+	Mg 2+	Sr 2+	Si 4+
	ºC		atm					mM
	25	7.2	10 -2.1	41	14.7	2.4	43.2	1.04	8.5	5.4	0.21	0.18
	50	6.86	10 -1.7	41	14	2.3	42	1.00	9.7	4	0.2	0.18
	70	6.66	10 -1.4	41	12	2.1	40	0.98	9.8	3	0.18	0.18
	80	6.57	10 -1.3	41	11	2.1	39	0.95	9.7	2.7	0.16	0.18
	90	6.5	10 -1.2	41	10	2.0	38	0.93	9.5	2.3	0.15	0.18

Table 3 . Chemical composition of the Callovo-Oxfordian porewater calculated by geochemical modelling accord- ing to the temperature increasing using the completed mineralogy. Temp pH pe pCO2 Cl - SO4 2+ TIC Na + K + Ca 2+ Mg 2+ Sr 2+ Fetotal Al 3+ Si 4+ ºC atm mM

 3

	25	7.2	-3	10 -2.1	41	15	2.4	43	1.0	8.5	5.4	0.21	47	2.9•10 -6 0.18
	50	6.66 -2.9 10 -1.3	41	14	3.8	42	1.0	9.9	4.1	0.2	0.05	2	0.35
	70	6.31 -2.8 10 -0.8	41	12	5.9	41	0.99	10	3.2	0.19	0.05	4	0.56
	80	6.15 -2.8 10 -0.5	41	11	7.6	40	0.98	10	2.8	0.18	0.05	6	0.69
	90	6.01 -2.8 10 -0.3	41	10	10	39	0.96	10	2.5	0.17	0.05	8	0.84

Table 4 . Summary of the literature analysing the stability constants of Ca-U(VI)-CO3 complexes. Data in grey have been superseded by later references of the same authors (2001). Reference log10 βº113 CaUO2(CO3)3 2- log10 βº213 Ca2UO2(CO3)3(aq) Ionic strength correction Technique

 4

	Bernard et al. (1996)	-	26.80±0.70 (I=0.1)	n.a.	TRLFS
	Kalmykov and Choppin (2000)	-	29.80 ± 0.70	SIT	fluorescence
	Bernhard et al. (2001)	25.40 ± 0.25	30.45 ± 0.35	Davies	TRLFS
	Dong and Brooks (2006)	27.18 ± 0.06	30.70 ± 0.05	Davies	Anion exchange resin
	Lee and Yun (2013)	27.27 ± 0.14	29.81 ± 0.19	SIT	TRLFS
					selective elec-
	Endrizzi and Rao (2014)	26.99 ± 0.72	30.84 ± 0.72	SIT	trode
					and UV/VIS
	Jo and Yun (2017)	26.13 ± 0.20	30.64 ± 0.20	SIT	TRLFS

Table 5 . Summary of the literature on the stability constants for Mg-U(VI)-CO3 complexes. Data in grey have been superseded by later references of the same authors.

 5

	Reference	log10 βº113 MgUO2(CO3)3 2-	log10 βº213 Mg2UO2(CO3)3(aq)	Ionic strength correction	Technique
	Dong and Brooks (2006)	26.11 ± 0.04		Davies	Anion exchange resin
	Dong and Brooks (2008)	25.02 ± 0.08		SIT	Anion exchange resin
	Dong and Brooks (2008)	25.80 ± 0.50		Davies	Anion exchange resin
	Geipel et al. (2008)	26.24 ± 0.13		Davies	TRLFS
	Endrizzi and Rao (2014)	26.25 ± 0.07		SIT	Selective electrode and UV/VIS
	Lee et al. (2017)	25.80 ± 0.30	27.10 ± 0.60	SIT	TRLFS

Table 6 . Stability constants for Ca-U(VI)-CO3 complexes at I = 0.1M and different temperatures suggested by Jo and Yun (2017).

 6

	TºC	log10 K (Reaction 1)	log10 K (Reaction 2)
	10 ± 0.1	2.46 ± 0.3	5.59 ± 0.7
	25 ± 0.2	3.01 ± 0.2	6.0 ± 0.3
	40 ± 0.5	2.86 ± 0.3	6.22 ± 0.3
	55 ± 0.6	3.17 ± 0.2	6.41 ± 0.6
	70 ± 1.0	-	6.72 ± 0.2

Table 7 . Summary of the batch experiments performed in the present work.

 7

	Type NaNO3	S/L	[Ca]	[U(VI)]initial Temperature	pH	CO2(g)
		M	g/L	nM	M	ºC		%
						20	7.2	1
	I	0.1	1 -25	0	≈ 5•10 -7	40 60	7.0 6.8	1 3
						80	6.6	5

Table 8 . Description of the tests performed in different temperatures and in closed system experiments.

 8

Table 10 . Estimated values for the log10 KX at different temperatures and pCO2 (atm).

 10

	Case Experimental conditions	log10 KX1	Estimated value ± corr. Coef.
	A	20 ºC; pCO2 = 10 -2.0 atm	1.08	0.06	0.995
	B	40 ºC; pCO2 = 10 -2.0 atm	1.19	0.06	0.999
	C	60 ºC; pCO2 = 10 -1.7 atm	1.53	0.06	0.998
	D	80 ºC; pCO2 = 10 -1.5 atm	1.68	0.16	0.966

Table 11 . Theoretical and measured pH, alkalinity and pCO2 of the batch experiments at different temperatures. Theor.: theoretical values according to the initial conditions of the experiments; Measured pH and alklinity in equilibrium; and calculated pCO2 in agreement with pH and alkalinity measured.

 11

	Temp. CO2(g) pH theor.	Alkalinity Theor.	pCO2 theor. pH measured	Alkalinity measured	pCO2 calculated
	20 ºC	1 %	7.2	2.0 meq/L	10 -2.1 atm	7.18	2.5 meq/L	10 -2.1 atm
	40 ºC	1 %	7.0	1.8 meq/L	10 -2.0 atm	7.16	2.5 meq/L	10 -2.0 atm
	60 ºC	3 %	6.8	1.8 meq/L	10 -1.5 atm	6.90	2.2 meq/L	10 -1.6 atm
	80 ºC	5 %	6.6	1.6 meq/L	10 -1.3 atm	6.87	2.2 meq/L	10 -1.5 atm

Table 12 . Results of the fitting for the stability constants of Ca-U(VI)-CO3 complexes using PEST.

 12

		20 ºC		40 ºC		60 ºC		80 ºC
	Estimated	value	±	value	±	value	±	value	±
	log10 β213

29.81 1.08 29.90 0.73 29.64 0.09 29.73 0.10 log10 β113 27.43 0.11 27.23 0.12 26.68 0.03 26.49 0.06

Table 13 . Stability constants of Ca-U(VI)-CO3 complexes calculated by 2D mapping method.

 13

		20 ºC	40 ºC	60 ºC	80 ºC
	Estimated	value	value	value	value
	log10 β213	29.35	29.9	29.6	29.72
	log10 β113	27.41	26.99	26.72	26.41

Table 14 . Thermodynamic data of Ca-U(VI)-CO3 ternary complexes determined in this work. Reaction log10 βº ± σ ΔrHº ± σ (kJ/mol) ΔrSº ± σ (kJ/K•mol) ΔrGº ± σ (kJ/mol)

 14

	Ca 2+ + UO 2 2+ +3CO 3 2-⇌ CaUO 2 (CO 3 ) 3 2-	27.3 ± 0.2 -27.4 ± 8.2 0.41 ± 0.004 -150.8 ± 8.2
	2Ca 2+ + UO 2 2+ +3CO 3 2-⇌ Ca 2 UO 2 (CO 3 ) 3 (aq)	29.7 ± 0.3	1.3 ± 1.8	0.59 ± 0.03	-177.1 ± 13

Table 15 . Calculated UO2•2H2O(am) solubility and aqueous U(VI) speciation under Callovo-Oxfordian conditions (Tournassat et al., 2008, Table 3)

 15

			Ca-U(VI)-CO3 stability constants
			Dong and Brooks (2006)	Present work
	Solid phase		UO2•2H2O(am)	UO2•2H2O(am)
	U concentration		1.9•10 -6 M	7.7•10 -7 M
		CaUO2(CO3)3 2-	24 %	18 %
	Speciation	Ca2UO2(CO3)3(aq)	75 %	78 %
		UO2(CO3)3 4-	1 %	2 %

Table 16 . Comparison of the thermodynamic parameters of the

 16

	Reaction	log10 βº ±σ	ΔrHº ±σ (kJ/mol)	ΔrSº ±σ (kJ/K•mol)	ΔrGº ±σ (kJ/mol)	Ref.
	Ca 2+ + UO 2 2+ +3CO 3 2-⇌CaUO 2 (CO 3 ) 3 2-	27.3 ±0.2 27.0 ±0.04	-27.4 ±8.2 0.43 ±0.01 -47 ±6 0.36 ±0.01	-155.8 ±8.2 p.w. -154.1 ±0.2 (1)
		29.7 ±0.3	0 ± 1.8	0.57 ±0.02	-169.5 ±1.8 p.w.
	2Ca 2+ + UO 2 2+ +3CO 3 2-⇌ Ca 2 UO 2 (CO 3 ) 3 (aq)	30.84 ±0.04	-47 ±7	0.43 ±0.01	-176.0 ±0.2	(1)
		30.75 ±0.2	-37 ±18	n.d.	n.d.	(2)

Ca-U(VI)-CO3 ternary complexes reactions. p.w. present work; (1) Endrizzi and Rao (2017); (2) Jo and Yun (2017); n.d. not defined.

  Table 17, we show some examples of different types of clay minerals, highlighting the most relevant for the present work.

Table 17 . Classification of the planar hydrated phyllosilicates and some examples for each clay mineral species (Brigatti et al., 2006). (the higlighted clay minerals are the ones relevant for the present study)

 17

	Layer type	Interlayer composition	Net layer charge	Group	Octahedral character	Example
					Tri	Amesite
	1:1	None or H2O	≈ 0	Serpentine-Kaolin	Di	Kaolinite
					Di-Tri	Odinite
		None	≈ 0	Talc-pyrophyllite	Tri Di	Talc Pyrophyllite
		Hydrated exchangeable cation	0.2-0.6	Smectite	Tri Di	Saponite Montmorillo-nite
		Hydrated exchangeable cation	0.6-0.9	Vermiculite	Tri Di	Vermiculite Vermiculite
	2:1	Exchangeable monova-lent cation	0.85-1.0 Micas	Tri Di	Phlogopite Muscovite
		Exchangeable and di-valent cation mono-	0.6-0.85 Illite	Tri Di	Biotite Illite
					Tri	Clinochlore
		Hydroxide sheet	variable	Chlorite	Di	Donbassite
					Di-Tri	Cookeite

Table 18 . Sorption parameters for U(VI) onto clay minerals of Callovo-Oxfordian formation (Bradbury and Bae- yens, 1997, 2004).

 18

		Montmorillonite	Illite
	CEC (eq/kg)	0.87	0.25

Specific surface area (m 2 /g) 39 20 Sites capacities (mol/kg)

	≡S s OH	2•10 -3	2•10 -3
	≡S w1,w2 OH	4•10 -2	4•10 -2

Table 19 . Cationic exchange and surface complexation reactions used in the sorption model.

 19

	Illite	Montmorillonite

Table 20 . Distribution of surface complexation reactions of U(VI) on illite and montmorillonite in Callovo-Oxfordian formation system (pH = 7.2 and pCO2 = 10 -2.1 atm) Surface complexation Illite Montmorillonite

 20

	Hydrolysis reactions	4.1 %	3.1 %
	Carbonato reactions	-	91.3 %

Table 21 . Description of the experimental cases performed on different phases of Callovo-Oxfordian formation at 20 ºC.

 21

	Case Methods	Material	Observations
		Sorption pH edge experiments in the		Validate the parameters describing the surface
	A	absence and in the presence of car-	Illite	complexation of hydrolysis and carbonate sur-
		bonates		face species
	B	Sorption isotherm experiments in the absence of carbonates	Illite and COx clay fraction	Sorption capacity of clay minerals for U-hy-drolysis and U-carbonate surface species
	C	Sorption isotherms experiments in COx porewater conditions	Illite, COx clay frac-tion and claystone	Impact of the formation of Ca-U(VI)-CO3 ter-nary complexes on the adsorption of U(VI) be-haviour

Table 22 . Summary of the experimental conditions for sorption of uranium in different components of Callovo- Oxfordian formation.

 22

	Experimental conditions	Sorption edges		Sorption isotherms
	Clayey material	Illite		Illite clay fraction	Illite clay fraction	Illite clay fraction	Illite claystone clay fraction	illite
	pCO2	0%	0.03%	0%	0.03%	1%	1%	5%
	pH range	4-10	4-10	7.2	7.3	7.3	7.3	6.6
	Ionic strength	0.1 (M) 0.1 (M)	0.1 (M)	0.1 (M)	0.1 (M)	0.1 (M)	0.1 (M)
	Background media	NaNO3	NaCl	NaNO3	NaCl	NaNO3	COx pw 3	NaNO3
	S/L (g/L)	2.0 0.5-1.0	2	2	2 2, 10 and 40	2.0

Table 23 . Alkalinity and pH measurements before and after the equilibrium with illite. pCO2 calculated from the equilibrium pH/alkalinity.

 23

	Alkalinity (meq/L)	pHinitial	pHfinal	pCO2 calculated i
	0	3.9	4.9
	0.036 (± 0.02)	6.1	6.0 6.4	-2.58; -3.10 -3.0; -3.53
	0.1 (± 0.05)	7.2	6.8 7.1	-2.96; -3.44 -3.26; -3.74
	0.96 (± 0.5)	8.0	8.1 8.2	-3.27;-3.78 -3.36; -3.86

Table 24 . Uranyl carbonate surface complexation reactions proposed for illite in Model III.

 24

	Carbonato surface complexation reactions	log10 K
	≡S s OH+UO 2 2+ + CO 3 2-⇌≡S s OUO 2 CO 3 -+2H +	9.8 i
	≡S s OH+UO 2 2+ + 2CO 3 2-⇌≡S s OUO 2 (CO 3 ) 2 3-+H +	17.5 (±0.5)
	≡S w1 OH+UO 2 2+ + CO 3 2-⇌≡S w1 OUO 2 CO 3 -+2H +	9.3 6

Table 25 . pH and alkalinity measurements after equilibrium with pCO2 atmospheric, 1% and 5% CO2(g) CO2(g) Measured pH Measured Alkalinity meq/L Air

 25

7.3 (± 0.09) 0.24 (± 0.09) 1%

7.3 (± 0.08) 2.5 (± 0.1) 5%

6.6 (± 0.01) 2.3 (± 0.1)

Table 26 . Surface species of U(VI) distribution on illite and montmorillonite for Callovo-Oxfordian formation ex- perimental conditions.

 26

		Total	main species	%
	Aqueous species	61 %	CaUO2(CO3)3 2 -Ca2UO2(CO3)3(aq)	81 % 17 %
	Surface complexes	39 %	≡S illite s ≡S mont OUO 2 (CO 3 ) 2 3-w1 OUO 2 CO 3 -	33 % 4 %

Table 28 . Compilation of the surface reactions for both illite and montmorillonite taken in for the model prediction. (n.a. not applied)

 28

	Illite	Montmorillonite

Table 29 . Description of the U(VI) sorption experiments carried out on illite, Callovo-Oxfordian clay fraction and clays samples as a function of temperature.

 29

	Case	Methods	Observations
		Sorption edges and isotherms of U(VI) in	Evaluate the correlation between aqueous and
		the absence of carbonates at different tem-	surface hydrolysis reactions on the surface of
		peratures	illite
		Sorption edges and isotherms of U(VI) in	Evaluate the correlation between aqueous and
	Illite	the presence of carbonates at different tem-	surface carbonate complexation reactions on
		peratures	the surface of illite
		Sorption isotherms of U(VI) in synthetic COx porewater composition at 80 ºC	Evaluate the influence of ternary complexes on the retention of U(VI) with the increase of tem-perature
		Sorption isotherms of U(VI) in the absence	Can we apply the results obtained for illite di-
	COx clay fraction	or presence of carbonates at 80 ºC Sorption isotherms of U(VI) in synthetic	rectly to smectite? Evaluate the influence of ternary complexes on
		COx porewater composition at 80 ºC	the retention with the increase of temperature
	COx claystone	Sorption isotherms of U(VI) in synthetic COx porewater composition at 80 ºC	Blind prediction in temperature

Table 30 . Experimental conditions of the sorption experimental cases for the different phases of Callovo-Oxfor- dian formation at different temperatures. (pw = porewater).

 30

	Experimental conditions	Sorption edges		Sorption isotherms
	Clayey material	illite		Illite clay fraction	Illite	Illite clay fraction	Illite clay fraction claystone
	Atmosphere	Ar(g)	air	Ar(g)	Air	1%	5%	5%
	Temperature	80 ºC	80 ºC	80 ºC	60-80 ºC	60-80 ºC	60-80 ºC	80 ºC
	pH range	4-10	4-10	Equilibrium	7	7.2	6.6	6.6
	Ionic strength (M)	0.1	0.1	0.1	0.1	0.1	0.1	0.1
	Background media	NaNO3	NaCl	NaNO3	NaCl	NaNO3	NaNO3	COx pw

Table 31 . Aqueous speciation of U(VI) (Guillaumont et al., 2003) and surface complexation reactions and respec- tive stability constants (Bradbury and Baeyens, 2009; Marques Fernandes et al., 2015) and enthalpy values (un- known: to be determined) used on the predictive model. (n.a.: not applied).

 31

	AQUEOUS COMPLEXATION		log10 Kº	ΔrHº kJ/mol
	UO 2 2+ + H 2 O⇌UO 2 (OH) + +H +		-5.25 (±0.24)	43.46 (±14.99)
	UO 2 2+ + 2H 2 O⇌UO 2 (OH) 2 +2H +		-12.15 (±0.07)	111.16 (±1.51)
	UO 2 2+ + 3H 2 O⇌UO 2 (OH) 3 -+3H +		-20.25 (±0.42)	148.06 i
	UO 2 2+ + 4H 2 O⇌UO 2 (OH) 4 2-+4H +		-32.4 (±0.68)	156.14 1
	UO 2 2+ + CO 3 2-⇌UO 2 CO 3		9.94 (±0.03)	5 (±2)
	UO 2 2+ + 2CO 3 2-⇌UO 2 (CO 3 ) 2 2-		16.61 (±0.09)	18.5 (±4)
	UO 2 2+ + 3CO 3 2-⇌UO 2 (CO 3 ) 3 4-		21.84 (±0.04)	-39.2 (±4.1)
			Illite	Montmorillonite
	SURFACE COMPLEXATION REACTIONS	log10 Kx	ΔrH kJ/mol	log10 Kx	ΔrH kJ/mol
	≡S s OH+UO 2 2+ ⇌≡S s OUO 2 + +H +	2.0	unknown	3.1	unknown
	≡S w1 OH+UO 2 2+ ⇌≡S w1 OUO 2 + +H +	0.1	unknown	0.7	unknown
	≡S s OH+UO 2 2+ + H 2 O⇌≡S s OUO 2 OH+2H +	-3.5	unknown	-3.4	unknown
	≡S w1 OH+UO 2 2+ + H 2 O⇌≡S w1 OUO 2 OH+2H +	-5.3	unknown	-5.7	unknown
	≡S s OH+UO 2 2+ + 2H 2 O⇌≡S s OUO 2 (OH) 2 -+3H +	-10.6	unknown	-11	unknown
	≡S s OH+UO 2 2+ + 3H 2 O⇌≡S s OUO 2 (OH) 3 2-+4H +	-19	unknown	-20.5	unknown
	≡S s OH+UO 2 2+ + CO 3 2-⇌≡S s OUO 2 CO 3 -+H +	n.a.	n.a.	9.8	unknown
	≡S w1 OH+UO 2 2+ + CO 3 2-⇌≡S w1 OUO 2 CO 3 -+H +	n.a.	n.a.	9.3	unknown
	≡S s OH+UO 2 2+ + 2CO 3 2-⇌≡S s OUO 2 (CO 3 ) 2 3-+H +	17.5	unknown	15	unknown

Table 33 . Surface complexation constants on strong sites together with the corresponding hydrolysis constants and enthalpy values for illite considered in Model IV. Surface species log10 OH Kx ΔrHº kJ/mol Aqueous species log10 OH K ΔrHº kJ/mol

 33

	≡S s OUO 2 +	2.0	43.46	UO 2 (OH) +	-5.25	43.46
	≡S s OUO 2 OH	-3.5	111.16	UO 2 (OH) 2 0	-12.15	111.16
	≡S s OUO 2 (OH) 2 -	-10.6	148.06	UO 2 (OH) 3 -	-20.25	148.06
	≡S s OUO 2 (OH) 3 2-	-19	156.14	UO 2 (OH) 4 2-	-32.4	156.14

Table 34 . Log10 OH Kxº fitted based on sorption edges and isotherms experiments performed at different tempera- tures for illite.

 34

	Log10Kxº	Sorption edges	Sorption Isotherms
	Surface species	20 ºC	80 ºC	20 ºC	60 ºC	80 ºC
	≡S s OUO 2 +	1.3	1.7	1.3	1.9	1.7
	≡S s OUO 2 OH	-3.3	-2.1	-3.3	-3.0	-2.1
	≡S s OUO 2 (OH) 2 -	-10.4	-7.2	-10.1	-8.6	-7.3
	≡S s OUO 2 (OH) 3 2-	-19.7	-14.3	-19.7	-16.5	-15.0

Table 36 . Alkalinity and pH measurements before and after the equilibrium with illite at 80 ºC. pCO2 calculated from the equilibrium pH/alkalinity.

 36

	pHinitial 20 ºC	Alkalinityinitial meq/L	pHfinal 80 ºC	Alkalinityfinal meq/L	pHfinal 20 ºC	pCO2 calculated i
	3.9	0	5.0	0
	6.1	0.036 (± 0.02)	6.5	0.036 ii (± 0.02)		-2.82/-3.39
			6.8			-2.18/-2.55
	7.4	0.17 (± 0.07)	7.2	0.25 (± 0.1)	7.7	-2.69/-3.07
			7.4			-2.92/-3.31
	8.5	2.5 (± 0.5)	8.2 8.2	2.5 (± 1.0)	8.8	-2.87/3.25 -2.89/-3.27
	9.31	18.1 (± 7.3)	9.2	18.5 (± 7.4)	9.3	-3.54/-3.94

Table 37 . pH of equilibrium and alkalinity measurements in the presence of carbonates at 60 and 80ºC.

 37

	Atmosphere	Temperature ºC	pH	Alkalinity (meq/L)	pCO2 calculated i
	Ar(g)	60 80	6.4 (±0.1) 6.5 (±0.2)	0.08 (±0.03) 0.08 (±0.03)	-2.5/-2.8 -2.5/-2.9
	Air	60 80	7.5 (±0.1) 7.5 (±0.1)	0.84(±0.3) 0.2(±0.1)	-2.6/-2.9 -3.1/-3.5
	1% CO2(g)	60 80	7.5 (±0.03) 7.6 (±0.03)	2.5(±1.0) 2.4(±0.9)	-2.0/-2.4 -2.2/-2.6
	5% CO2(g)	60 80	6.8 (±0.01) 7.0 (±0.02)	2.5(±1)	-1.4/-1.7

Table 38 . pH and alkalinity measurements in equilibrium for the sorption isotherms on clay fraction at 80 ºC in different atmospheres. Atmosphere pHfinal Alkalinity (meq/L) pCO2 calculated i

 38

Table 39 . Enthalpy values for the hydrolysis surface complexation reactions for illite and montmorillonite consid- ered in the present study.

 39

	SURFACE COMPLEXATION REACTIONS	log10 K	Illite ΔrH kJ/mol log10 K Montmorillonite ΔrH kJ/mol
	≡S s OH+UO 2 2+ ⇌ ≡S s OUO 2 + +H +	2.0	-10	-10
	≡S w1 OH+UO 2 2+ ⇌ ≡S w1 OUO 2 + +H +	0.1	-10	-10
	≡S s OH+UO 2 2+ + H 2 O⇌ ≡S s OUO 2 OH+2H +	-3.5	48	48
	≡S w1 OH+UO 2 2+ + H 2 O⇌≡S w1 OUO 2 OH+2H +	-5.3	48	48
	≡S s OH+UO 2 2+ + 2H 2 O⇌ ≡S s OUO 2 (OH) 2 -+3H +	-10.6	121	121
	≡S s OH+UO 2 2+ + 3H 2 O⇌ ≡S s OUO 2 (OH) 3 2-+4H +	-19	156	156
	≡ 𝑆 𝑠 𝑂𝐻 + 𝑈𝑂 2 2+ + 𝐶𝑂 3 2-⇌≡ 𝑆 𝑠 𝑂𝑈𝑂 2 𝐶𝑂 3 -+ 𝐻 +	n.a.	n.a.	9.8	0
	≡S w1 OH+UO 2 2+ + CO 3 2-⇌≡S w1 OUO 2 CO 3 -+H +	n.a.	n.a.	9.3	0
	≡S s OH+UO 2 2+ + 2CO 3 2-⇌≡S s OUO 2 (CO 3 ) 2 3-+H +	17.5	0	15	0

Table 40 . Experimental pH and alkalinity measurements at 20 and 80 ºC for illite, Callovo-Oxfordian clay fraction and claystone samples in synthetic Callovo-Oxfordian porewater in equilibrium with 1% and 5% of CO2(g), respec- tively.

 40

		Temperature	Atmosphere	pH	Alkalinity (meq/L) pCO2 calculated i
	Illite	20 ºC 80 ºC	1% CO2(g) 7.1 (± 0.02) 5% CO2(g) 6.8 (± 0.01)	2.5 (± 0.1) 2.2 (± 0.8)	-2.06/-2.09 -1.4/-1.8

Table 41 . Aqueous and surface speciation at 20 and 80ºC for the Callovo-Oxfordian formation system in equilib- rium with synthetic porewater for [U(VI)] = 10 -7 M.

 41

			20 ºC				80 ºC
		Total main species		Total main species
	Aqueous species	61 % CaUO2(CO3)3 2-Ca2UO2(CO3)3(aq)	81% 17%	5 %	CaUO2(CO3)3 2-Ca2UO2(CO3)3(aq)	43% 47%
		39 % ≡S illite s	OUO 2 (CO 3 ) 2 3-	33% 95 % ≡S illite s	OUO 2 (CO 3 ) 2 3-	22%
	Surface complexes	≡S mont w1 OUO 2 CO 3 -	4%		Hydrolysis surface com-plexes of illite	64%

Table 42 . Proposed thermodynamic parameters for the surface complexation reactions considered in the model approach.

 42

		Illite	Montmorillonite
	log10 K	ΔrH kJ/mol	log10 K	ΔrH kJ/mol

  Table A. 3. The final chemical composition of the synthetic porewaters are provided in Table C. 7.

Table A . 2. Theoretical chemical composition of the Callovo-Oxfordian porewater at 20 ºC and 80ºC.

 A

	mM	Ca	Mg	Na	K	Cl	SO4 2-	Sr	CO3 2-
	20 ºC 8.55	5.42	43.35	1.04	40.8	14.75	0.21	2.5
	80 ºC 9.35	2.65	36.59	0.95	40.8	10.77	0.17	2.4

Table A . 3. Salt content in g/L used in preparation of synthetic COx porewater at 20 and 80ºC

 A

	Salt product	Purity (%) Provided	20 ºC (g/L) 80 ºC (g/L)
	SrCl2 • 6H2O	100%	Aldrich	0.0562	0.0453
	KCl	99%	Fisher	0.0786	0.0715
	MgCl2 • 6H2O	99%	Merck	1.1127	0.5442
	CaCl2 • 2H2O	99%	Aldrich	1.2697	1.3885
	NaCl	99.5%	Sigma Aldrich	0.6666	0.8839
	Na2SO4	99%	Panréac	2.1163	3.5051
	NaHCO3	99.5%	Sigma Aldrich	0.2111	0.2025
	A.6. Solid phases:
	Anion Exchange resin:

Table A . 4. Mineral composition of Callovo-Oxfordian claystone

 A

	Mineral	Wt%
	Illite	17
	IS R0	28

Table A . 5. Experimental conditions for the batch experiments. Temp.= temperature; COx = Callovo-Oxfordian; I = Ionic strength.

 A

	A.7.1.
	[U(VI)]		µM			0.5		0.05		0.5			0.5		0.1	0.03			0.05-20				0.01-10	0.05-20
	[Ca]		mM			0			0-5					0			0		0		8.5		0	8.5	0	8.5	8.5
	[Carb.]		mM		0.24	2.0	0.24		2.0			0		0	0.24	0	0.24-2	2.0	2.0	0-2.0	2.0	0-2.0	2.0	2.01
	I		M				0.1								0.1			0.1			0.1				0.1	0.1
		Background				NaNO3	NaNO3 or	NaClO4		NaNO3		NaCl			COx pw	NaNO3	NaCl	NaNO3	NaCl and-	NaNO3	NaNO3	NaNO3	NaNO3	COx porewater	NaNO3	COx porewater	COx porewater
	Range of S/L		g/L			1-25			2.5				0.5			50	2.0	0.5-1.0		2		10	2		2	40	5
		pH			8.1	7.2-6.6	8.1		7.2-6.6		equilibrium			7.2		4-10	equilibrium	7.2, 6.6 6.8 and	7.2	6.6	7.2			6.6	7.2	6.6
	Time		days			5			5					3, 7 1, 2,	14 and		3			3				3	3
	pCO2		atm		atmospheric	1, 3 and 5%	atmospheric		1, 3 and 5%		atmospheric			1%	inert	atmospheric	inert	atmospheric,	1% and 5%	1%	5%	Inert 1% and	1%	Inert 5% and	5%	1%	5%
	Temp.	Exp.	ºC	Anionic-exchange resin	20	20, 40, 60 ΔS/L	and 80	20	ΔCa	20, 40, 60	and 80	Kinetics	illite	20	COx claystone	Sorption edge	20 and 80 illite	Sorption isotherms	80 20, 60 and	illite	20	80	20	COx clay	fraction	80	20 COx claystone 80

Open system (atmospheric conditions) procedure:

Table C . 1. Experimental test of the control of the temperature using an isolation-aluminium box without external layer of Teflon. Open-box closed box measure by a small aperture closed box pre-heated at 90ºC 1 st test duplicated 2 nd test 3 rd test Time Temp Time Temp Time Temp corrected Temp time Temp corrected Temp

 C

		Open-box		closed box measure by a small aperture	closed box pre-heated at 90ºC
		1 st test	duplicated			2 nd test	3 rd test
	Time Temp Time Temp Time Temp corrected Temp	time	Temp	corrected Temp
	min	°C	min	°C	min	°C	°C	min	°C	°C
		68.4	12	66.8	19	70	73	19	70	75
		67.9	13	65.9	21.25 69	72	21.25 69	74
		67.1	14	65	23.5	68	71	23.5	68	73
		66.3	15		25.5	67	70	25.5	67	72
		65.6	16	64.1	28	66	69	28	66	71
		64.6	17	62.8	30	66	69	30	66	71
	min	°C	min	°C	min	°C	°C	min	°C	°C
	0	76	0	77.8	0	74	77	0	74	79
	1	75.9	1	77.1	1	74	77	1	74	79
	2	75.8	2	76.1	1.5	75	78	1.5	75	80
	3	75.4	3		3	75	78	3	75	80
	4	74.7	4	74.3	4	75	78	4	75	80
	5	74.1	5	73.4	5	75	78	5	75	80
	6	73.3	6	72.4	7	75	78	7	75	80
	7	72.5	7	71.4	10	75	78	10	75	80
	8	71.9	8	70.4	10.5	74	77	10.5	74	79
	9	71	9	69.5	12.66 73	76	12.66 73	78
	10	70.1	10	68.6	15	72	75	15	72	77

Table C . 2. Experimental test of the control of the temperature using an isolation-aluminium box with external layer of Teflon. pre-heating at 90 ºC special closing system (pre-heated 90 ºC) 4 th test 5 th test time Temp corrected Temp time Temp corrected Temp

 C

		pre-heating at 90 ºC	special closing system (pre-heated 90 ºC)
		4 th test		5 th test
	time	Temp	corrected Temp	time	Temp	corrected Temp
	min	°C	°C	min	°C	°C
	20	75	77	13	81	84
	22	74	76	14	81	84
	24.33	73	75	15	81	84
	27.2	72	74	16	80	83
	29.75	71	73	17	80	83
	min	°C	°C	min	°C	°C
	0.5	77	79	0	76	79
	2	77	79	1	77	80
	3.33	78	80	2	78	81
	5	78	80	3	78	81
	7	78	80	4	79	82
	9	78	80	5	80	83
	10	78	80	6	80	83
	11.33	77	79	7	80	83
	13	77	79	8	80	83
	15	77	79	9	81	84
	15.5	76	78	10	81	84
	17	76	78	11	81	84
	18.66	75	77	12	81	84

Table C . 3. Experimental test for sealing control.

 C

	Test	Control	Observations
	0	Fe(II)/Fe(III)	Closed tubes and agitated in open atmosphere
		pH
	1	Fe(II)/Fe(III)	Closed tubes and agitated inside of the glove-box for 7 days
		pH

Table C . 5. Uncertainty of the signal of the TOC Analyser determined using standard solution of IC, TC and ultra- pure water.

 C

		Ultra-pure water	Standard TC		Standard IC		Uncertainty
	Date	TC	IC	theoretical	measured	±	theoretical	measured	±	TC	IC
		ppm	ppm	ppm	ppm	ppm	ppm	ppm	ppm	%	%
	11/06/15 -0.035	-0.004	5	5.09		5	5.18		2.4	3.7
	25/06/15 -0.097	-0.019	2	1.837		2	1.838		3.3	7.2
	08/09/15	0.015	-0.022	2	1.958	0.13	2	1.977	0.050	2.9	0.05
	08/09/15	/	/	2	2.155	0.07	/	/	/	7.0	/
	10/09/15 -0.058	-0.004	2	1.949	0.09	2	1.862	0.008	0.4	6.7
	10/09/15	/	/	2	1.992	0.03	2	2.073	0.07	2.5	3.9
	25/09/15	0.010	-0.017	2	1.959	0.12	2	1.921	0.02	2.6	3.1
	25/09/15	0.144	0.024	2	2.089	0.05	2	2.146	0.02	4.0	8.2

Table C . 6. TOC measurements of the experimental tests.

 C

	ID ID Test_ tube_COT_60°C_E2	TC (ppm) TC (ppm) valeur ± (k=2) valeur ± (k=2) 0.81 0.03	IC(ppm) IC(ppm) valeur ± (k=2) valeur ± (k=2) 0.1930 0.0157	TOC éch. (ppm) TOC éch. (ppm) valeur ± (k=2) valeur ± (k=2) 0.62 0.04
	Test_ blanc_eau UP 25/06 Test_ tube_COT_60°C_E3	0.13 0.45	0.01 0.02	0.0420 0.0690	0.0031 0.0056	0.084 0.38	0.007 0.02
	Test_ tube_1_80°C_COT_t1 Test_ tube_COT_80°C_E1	2.42 2.20	0.12 0.09	0.0040 0.5710	0.0003 0.0465	2.413 1.63	0.116 0.10
	Test_ tube_1_80°C_COT_t2 Test_ tube_COT_80°C_E2	0.94 1.70	0.07	/	0.4880 0.3420	/ 0.0279	0.449 1.36	0.07	/
	Test_ tube_1_80°C_COT_t3 Test_ tube_COT_80°C_E3 Test_ tube_1_80°C_COT_t4 Test_ tube_blanc_22/09 Test_ tube_1_80°C_COT_t5 Test_ tube_COT_20°C_E1	1.33 0.71 1.15 17.32 0.75 -0.01	0.03 / 0.68 / 0.00	/	0.4630 0.0870 0.7440 3.5300 0.2540 0.0170	/ 0.0071 / 0.2877 / 0.0014	0.862 0.62 0.409 13.79 0.500 -0.03	0.03 0.74 / / 0.00	/
	Test_ tube_1_80°C_COT_t6	1.39	/		0.9290	/	0.464	/
	Test_ tube_1_80°C_COT_t7 Test_ tube_2_80°C_COT_t1 Test_ tube_2_80°C_COT_t2 C.2. Chemical composition of stock solutions: 1.29 / 0.5030 / 0.788 / 1.60 / 0.6580 / 0.939 / 1.27 / 0.3260 / 0.944 /
	Test_ tube_2_80°C_COT_t3	1.35	/		0.8510	/	0.498	/
	Test_ tube_2_80°C_COT_t4	1.48	/		0.7110	/	0.764	/
	Test_ tube_2_80°C_COT_t5	1.19	/		0.2470	/	0.945	/
	Test_ tube_1_20°C_COT_t1	0.74	/		0.6230	/	0.120	/
	Test_ tube_1_20°C_COT_t2	1.17	/		0.9080	/	0.260	/
	Test_ tube_1_20°C_COT_t3	0.64	/		0.5750	/	0.064	/
	Test_ tube_blanc 10/09	0.40	0.01		0.4760	0.0319	-0.08	0.03
	Test_ tube_COT_20°C_A0	1.15	0.08		0.8410	0.0004	0.31	0.08
	Test_ tube_COT_20°C_A1	0.20	0.01		0.2400	0.0001	<lq
	Test_ tube_COT_20°C_B0	1.12	0.08		0.7130	0.0004	0.40	0.08
	Test_ tube_COT_20°C_B1	0.26	0.02		0.2040	0.0001	0.05	0.02
	Test_ tube_COT_20°C_B2	0.17	0.01		0.1490	0.0001	0.022	0.012
	Test_ tube_COT_20°C_C1	1.22	0.09		0.7650	0.0004	0.46	0.09
	Test_ tube_COT_20°C_C2	0.17	0.01		0.1680	0.0001	<lq
	Test_ tube_COT_20°C_D0	0.77	0.05		0.7270	0.0004	0.04	0.05
	Test_ tube_COT_60°C_A0	1.91	0.13		0.7060	0.0004	1.20	0.13
	Test_ tube_COT_60°C_A1	0.88	0.06		0.2410	0.0001	0.64	0.06
	Test_ tube_COT_60°C_A2	0.52	0.04		0.1390	0.0001	0.38	0.04
	Test_ tube_COT_60°C_B0	1.78	0.12		0.5770	0.0003	1.20	0.12
	Test_ tube_COT_60°C_B1	0.52	0.04		0.1660	0.0001	0.36	0.04
	Test_ tube_COT_60°C_B2	0.53	0.04		0.2450	0.0001	0.29	0.04
	Test_ tube_COT_60°C_C1	1.40	0.10		0.6700	0.0003	0.73	0.10
	Test_ tube_COT_60°C_C2	0.54	0.04		0.2370	0.0001	0.31	0.04
	Test_ tube_COT_60°C_C3	0.46	0.03		0.2320	0.0001	0.23	0.03
	Test_ tube_COT_60°C_D0	1.04	0.07		0.6710	0.0003	0.37	0.07
	Test_ tube_COT_60°C_D1	0.46	0.03		0.2050	0.0001	0.26	0.03
	Test_ tube_COT_60°C_D2	0.38	0.03		0.2090	0.0001	0.17	0.03
	Test_ tube_COT_80°C_A0	2.84	0.20		0.8260	0.0004	2.0	0.2
	Test_ tube_COT_80°C_A1	0.94	0.07		0.1590	0.0001	0.78	0.07
	Test_ tube_COT_80°C_A2	1.05	0.03		0.2440	0.0163	0.81	0.03
	Test_ tube_COT_80°C_B0	2.87	0.07		0.7010	0.0470	2.17	0.09
	Test_ tube_COT_80°C_B1	1.51	0.04		0.4690	0.0314	1.04	0.05
	Test_ tube_COT_80°C_B2	1.03	0.03		0.1850	0.0124	0.84	0.03
	Test_ tube_COT_80°C_C1	1.91	0.05		0.1840	0.0123	1.72	0.05
	Test_ tube_COT_80°C_C2	1.25	0.03		0.2420	0.0162	1.00	0.04
	Test_ tube_COT_80°C_C3	0.99	0.02		0.1830	0.0123	0.80	0.03
	Test_ tube_COT_80°C_D0	1.24	0.03		0.0720	0.0048	1.17	0.03
	Test_ tube_COT_80°C_D1	0.72	0.02		0.2280	0.0153	0.49	0.02
	Test_ tube_COT_80°C_D2	0.61	0.02		0.1970	0.0132	0.41	0.02
	Test_ tube_blanc_22/09	0.58	0.02		0.6160	0.0502	<lq	0.06
	Test_ tube_COT_20°C_E1	0.89	0.03		0.6430	0.0524	0.24	0.06
	Test_ tube_COT_20°C_E2	0.29	0.01		0.3140	0.0256	<lq	0.03
	Test_ tube_COT_60°C_E1	1.18	0.05		0.5640	0.0460	0.62	0.07

Table C . 7. Chemical composition of the synthetic Callovo-Oxfordian porewater. Uncertainty of k=2

 C

	Porewater	Na	Ca	K	Mg	Cl	SO42+
	1% CO2(g) 42.8(±1.3) 8.3(±0.3) 1.06(±0.16) 5.4(±0.2) 40.2(±1.4) 13.9(±0.6)
	5%CO2(g) 42.9(±1.3) 8.4(±0.3) 0.91(±0.17) 5.4(±0.2)

Table C . 8. Measurements of the pH for the sorption edge experiments in atmospheric pCO2. (1) used only in the experiments at 80ºC; (2) this solution was not taken for the experiments performed at 20ºC.

 C

		bubbling	After overnight agitation
		Ar/CO2(g)	in pCO2 atmospheric
	pH Theoretical	time	pH final
	3.8	-	3.9
	6	-	6.1
	7	1h	7.2
	7.5	30 min	7.4 (1)
	8	1h	8.0
	8.5	50 min	8.5 (1)
	9.3	3h	9.8 (2)
	C.3

. Speciation of Ca-U(VI)-CO 3 species experi- mental results: Table C. 9. Alkalinity measurements using H2SO4 (0.16 N) in the blank samples of the batch experiments using anionic-exchange resin. *measured

	Experiment	T	NaHCO3 (0.01M) pH	digits acid	vol. sample	Alkalinity meq/L
	E1A E1B	20ºC	600 µL	7.43 7.18	105 124	10 10	2.10 (±0.1) 2.48 (±0.12)
	E2A E2B	40ºC	595 µL	7.25* 7.17	115.5 124	10 10	2.16 (±0.11) 2.48 (±0.12)
	E3A E3B	60ºC	580 µL	6.88 6.93	114 100	10 10	1.90 (±0.1) 2.23 (±0.1)
	E4A	80ºC	568 µL	6.91	109	10	2.18 (±0.1)

Table C . 13. Experimental results of the ionic exchange method at 40 ºC in closed system at 1% of CO2(g) in the absence of calcium and varying the solid-liquid ration, in 0.1M of NaNO3 adding initially

 C

	resin		S/L	pH	T	U-238 in solution	U sorbed		KD
	ID			final
	g	±	g/L	±	ºC	mol/L	±	mol/g	±	mL/g ±
	Mother sol.

[U(VI)] = 5•10 -7 M. (Ba: blank sample without U(VI); Bb: blank sample without resin with U)

Table C . 14. Experimental results of the ionic exchange method at 40 ºC in closed system at 1% of CO2(g) in the presence of calcium, 0.1 M of NaNO3 at fixed solid-liquid ratio adding initially [U(VI)] = 5•10 -7 M. (blank samples in the absence of calcium or presence of calcium for Ba: without U(VI); and Bb: without resin with U)

 C

	ID Mother sol. E2B-0 E2B-1 E2B-2 E2B-3 E2B-4 E2B-5 E2B-6 906 C.5. Illite: sorption experiments Ca resin pH T U-238 in solution U-sorbed KD Final ºC mM g ± mol/L ± mol/g ± mL/g / / / / 5.21E-05 1.26E-05 / / / 0 0.0637 0.002 7.16 38.8 5.17E-08 1.13E-08 1.84E-07 5.00E-08 3564 1244 ± / 0.1 0.0621 0.002 7.18 35 6.09E-08 1.30E-08 1.85E-07 5.13E-08 3040 1064 0.2 0.0627 0.002 7.19 38 7.61E-08 1.43E-08 1.77E-07 5.09E-08 2333 800 0.2 0.0622 0.002 7.18 39 8.70E-08 1.51E-08 1.74E-07 5.13E-08 2006 685 0.2 0.0633 0.002 7.16 38.3 9.79E-08 1.60E-08 1.67E-07 5.04E-08 1707 586 0.2 0.0622 0.002 7.15 37.9 1.08E-07 1.72E-08 1.66E-07 5.14E-08 1530 533 1 0.0627 0.002 7.17 38 1.59E-07 2.02E-08 1.44E-07 5.11E-08 C.5.1. Sorption isotherms:

Table C . 23. Alkalinity measurements in sorption isotherms experiments of U(VI)/illite system (2 g/L) at different temperatures. *synthetic Callovo-Oxfordian porewater.

 C

	atmosphere T	pH	Acid H2SO4 Vol. acid	±	Alkalinity	±
	ºC		N	digits		meq/L
	Ar(g)	6.57	0.02	31 13.0	0.08	0.03
	Atmospheric	7.3	0.02	94 37.6	0.24	0.09
	1% CO2(g)	7.26	0.2	127 50.8	2.54	1.02
	5% CO2(g)	6.62	0.2	114 45.6	2.28	0.91
	1% CO2(g)*	7.18	0.2	127 50.8	2.54	1.02
	Ar(g)	6.48	0.02	33 13.2	0.08	0.03
	Atmospheric	7.57	0.02	96	134	0.24	0.34
	1% CO2(g)	7.45	0.2	127 50.8	2.54	1.02
	5% CO2(g)	6.89	0.2	127 50.8	2.54	1.02
	Ar(g)	6.53	0.02	31 12.4	0.08	0.03
	Atmospheric	7.11	0.02	91 36.4	0.23	0.09
	1% CO2(g)	7.67	0.2	118 47.2	2.36	0.94
	5% CO2(g)	7.08	0.2	118 47.2	2.36	0.94
	5% CO2(g)*	6.85	0.2	108 43.2	2.16	0.86

Table C . 24. Experimental results of the sorption isotherms of U(VI) onto illite at 20 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.052	2.08	0.05	2.21E-05	1.20E-06	6.01	22.1	3.11E-08	2.06E-09	1.06E-05	6.35E-07	341978	30541
	6	0.051	2.04	0.05	8.29E-06	4.45E-07	6.28	22.1	1.32E-08	1.26E-09	4.06E-06	2.41E-07	306699	34423
	5	0.049	1.96	0.05	2.76E-06	1.43E-07	6.44	22.1	7.10E-09	3.36E-10	1.41E-06	8.17E-08	198137	14844
	4	0.054	2.16	0.05	5.54E-07	3.17E-08	6.52	22.2	6.64E-09	3.53E-10	2.53E-07	1.59E-08	38120	3138
	3	0.05	2	0.05	2.77E-07	1.57E-08	6.52	22.3	3.09E-09	1.72E-10	1.37E-07	8.59E-09	44316	3720
	2	0.053	2.12	0.05	1.29E-07	7.06E-09	6.52	22.3	1.43E-09	1.22E-10	6.03E-08	3.65E-09	42181	4412
	1	0.05	2	0.05	4.61E-08	2.70E-09	6.46	22.1	3.66E-10	1.05E-10	2.29E-08	1.47E-09	62603	18436
	9-BNF	/	/	/	2.21E-05	1.30E-06	4.76	22.2	2.50E-05	1.18E-06	/	/	/	/
	9-BF	/	/	/	2.21E-05	1.30E-06	4.76	22.2	2.52E-05	1.34E-06	/	/	/	/
	8-BNF	/	/	/	4.61E-08	2.64E-09	5.57	22.3	4.81E-08	2.27E-09	/	/	/	/
	8-BF	/	/	/	4.61E-08	2.64E-09	5.57	22.3	4.34E-08	1.93E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	illite	S/L		U(VI) ini		pH	Temp	U(VI) eq.		sorb. U(VI)		KD

Table C . 25. Experimental results of the sorption isotherms of U(VI) onto illite at 20 ºC in atmospheric conditions in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.05	2	0.05	1.73E-05	9.08E-07	7.11	20.7	2.56E-08	1.51E-09	8.40E-06	5.02E-07	17583	1476
	6	0.049	1.96	0.05	6.48E-06	3.36E-07	7.27	20.5	4.78E-07	2.82E-08	3.23E-06	1.91E-07	22032	1687
	5	0.049	1.96	0.05	2.16E-06	1.08E-07	7.33	20.7	1.47E-07	7.14E-09	1.04E-06	6.15E-08	8060	814
	4	0.052	2.08	0.05	3.17E-07	1.74E-08	7.34	20.6	1.29E-07	1.05E-08	1.48E-07	9.19E-09	15988	3348
	3	0.05	2	0.05	1.58E-07	8.60E-09	7.36	20.3	9.24E-09	1.85E-09	7.77E-08	4.74E-09	26035	4336
	2	0.05	2	0.05	7.39E-08	3.86E-09	7.35	20.7	2.98E-09	4.62E-10	3.61E-08	2.15E-09	23262	6438
	1	0.049	1.96	0.05	2.64E-08	1.49E-09	7.31	20.7	1.55E-09	4.20E-10	1.32E-08	8.37E-10	29717	11370
	9-BNF	/			1.73E-05	9.74E-07	6.76	20.3	4.45E-10	1.68E-10	/	/	/	/
	9-BF	/			1.73E-05	9.74E-07	6.76	20.3	1.80E-05	8.82E-07	/	/	/	/
	8-BNF	/			2.64E-08	1.45E-09	7.3	20.6	1.78E-05	7.98E-07	/	/	/	/
	8-BF	/			2.64E-08	1.45E-09	7.3	20.6	2.67E-08	1.26E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	illite	S/L		U(VI) ini		pH	T	U(VI) eq.		U(VI) sorb.		KD

Table C . 26. Experimental results of the sorption isotherms of U(VI) onto illite at 20 ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	9	0.053	2.12	0.05	1.73E-05	2.44E-06	7.29	21.1	9.03E-06	4.16E-07	3.89E-06	1.17E-06	430.5	131
	8	0.051	2.04	0.05	6.48E-06	9.13E-07	7.28	21.1	4.53E-06	2.14E-07	9.54E-07	4.60E-07	210.3	102.1
	7	0.05	2	0.05	1.69E-06	2.39E-07	7.09	22	7.66E-07	3.95E-08	4.64E-07	1.22E-07	606.5	161.9
	6	0.049	1.96	0.05	6.35E-07	8.94E-08	7.21	20.8	2.41E-07	1.13E-08	2.01E-07	4.63E-08	833.7	195.8
	5	0.051	2.04	0.05	2.12E-07	2.97E-08	7.28	21.2	5.46E-08	2.98E-09	7.71E-08	1.47E-08	1411.7	280.9
	4	0.05	2	0.05	1.23E-08	1.79E-09	7.33	21.1	1.56E-09	1.39E-10	6.28E-09	1.04E-09	4016.1	755.6
	3	0.05	2	0.05	6.15E-09	8.93E-10	7.32	20.9	7.35E-10	1.39E-10	3.16E-09	5.23E-10	4300	1079.1
	2	0.05	2	0.05	2.87E-09	4.15E-10	7.32	21	7.98E-10	1.26E-10	1.25E-09	2.48E-10	1563.2	397
	1	0.051	2.04	0.05	1.03E-09	1.50E-10	7.33	20.6	3.66E-10	1.18E-10	3.98E-10	1.03E-10	1087.4	448.5
	12-BNF	/	/	/	1.69E-06	2.47E-07	7.28	21.2	1.75E-05	8.40E-07	/	/	/	/
	12-BF	/	/	/	1.69E-06	2.47E-07	7.28	21.2	1.76E-05	7.98E-07	/	/	/	/
	11-BNF	/	/	/	1.69E-06	2.47E-07	7.09	21.7	1.67E-06	7.56E-08	/	/	/	/
	11-BF	/	/	/	1.69E-06	2.47E-07	7.09	21.7	1.67E-06	7.98E-08	/	/	/	/
	10-BNF	/	/	/	1.03E-09	1.49E-10	7.31	21.5	1.15E-09	1.39E-10	/	/	/	/
	10-BF	/	/	/	1.03E-09	1.49E-10	7.31	21.5	1.18E-09	1.30E-10	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini				eq.		sorb.
	ID	illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 27. Experimental results of the sorption isotherms of U(VI) onto illite at 20 ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.05	2	0.05	2.16E-05	1.20E-06	6.6	21.8	2.29E-06	1.22E-07	9.68E-06	6.50E-07	4235.3	363.1
	6	0.05	2	0.05	8.12E-06	4.43E-07	6.62	20.9	4.79E-07	3.66E-08	3.82E-06	2.43E-07	7981.1	793.3
	5	0.05	2	0.05	2.71E-06	1.43E-07	6.61	21	1.14E-06	5.46E-08	7.82E-07	7.91E-08	683.8	76.5
	4	0.05	2	0.05	3.19E-07	1.65E-08	6.61	21.1	6.59E-08	3.24E-09	1.30E-07	9.20E-09	1966.5	169.7
	3	0.049	1.96	0.05	1.59E-07	8.15E-09	6.63	21	2.25E-08	2.18E-09	7.15E-08	4.76E-09	3180.4	374.8
	2	0.052	2.08	0.05	7.44E-08	3.64E-09	6.61	21.3	7.27E-09	5.88E-10	3.30E-08	1.99E-09	4538.1	457.9
	1	0.05	2	0.05	2.66E-08	1.42E-09	6.6	21.2	2.15E-09	1.68E-10	1.25E-08	7.95E-10	5798.8	584.6
	9-BNF	/	/	/	2.16E-05	1.15E-06	6.61	21.7	2.15E-05	1.01E-06	/	/	/	/
	9-BF	/	/	/	2.16E-05	1.15E-06	6.61	21.7	2.13E-05	9.66E-07	/	/	/	/
	8-BNF	/	/	/	2.66E-08	1.38E-09	6.63	21.9	2.69E-08	1.30E-09	/	/	/	/
	8-BF	/	/	/	2.66E-08	1.38E-09	6.63	21.9	2.71E-08	1.26E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 28. Experimental results of the sorption isotherms of U(VI) onto illite at 20 ºC in 1% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.25	9.994	0.165	5.35E-08	2.87E-09	7.18		5.42E-08	3.61E-09	1.54E-09	3.54E-10	40.6	9.5
	6	0.249	9.96	0.164	1.73E-05	9.08E-07	7.08	20.9	1.42E-05	6.72E-07	3.13E-07	1.14E-07	22.1	8.1
	5	0.249	13.33	0.22	6.94E-06	3.73E-07	7.11	20.9	5.01E-06	2.31E-07	1.44E-07	3.30E-08	28.8	6.7
	4	0.249	9.956	0.164	2.16E-06	1.08E-07	7.13	20.9	1.59E-06	7.14E-08	5.69E-08	1.30E-08	35.7	8.3
	3	0.249	9.964	0.164	6.42E-07	3.44E-08	7.14	20.7	4.44E-07	2.02E-08	1.99E-08	4.06E-09	44.7	9.4
	2	0.251	10.028	0.165	2.14E-07	1.09E-08	7.14	20.6	1.41E-07	6.30E-09	7.27E-09	1.27E-09	51.7	9.3
	1	0.25	9.994	0.165	5.35E-08	2.95E-09	7.12	20.3	3.80E-08	1.89E-09	1.54E-09	3.54E-10	40.6	9.5
	9-BNF	/	/	/	1.73E-05	9.52E-07	7.1	21.6	1.71E-05	8.40E-07	/	/	/	/
	9-BF	/	/	/	1.73E-05	9.52E-07	7.1	21.6	1.75E-05	8.40E-07	/	/	/	/
	8-BNF	/	/	/	5.35E-08	2.87E-09	7.18	21.6	5.31E-08	3.82E-09	/	/	/	/
	8-BF	/	/	/	5.35E-08	2.87E-09	7.18	21.6	5.42E-08	3.61E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 29. Experimental results of the sorption isotherms of U(VI) onto illite at 60 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.05	2	0.05	2.21E-05	1.20E-06	6.38	61.4	3.63E-08	1.64E-09	1.10E-05	6.63E-07	304129.6	22856.9
	6	0.049	1.96	0.05	8.29E-06	4.45E-07	6.56	61.4	4.47E-08	2.02E-09	4.21E-06	2.52E-07	94057.2	7046.3
	5	0.05	2	0.05	2.76E-06	1.43E-07	6.64	62	2.35E-08	1.05E-09	1.37E-06	7.98E-08	58250	4272.8
	4	0.05	2	0.05	5.54E-07	3.17E-08	6.65	62.3	7.30E-09	3.57E-10	2.73E-07	1.73E-08	37427.5	2997.5
	3	0.053	2.12	0.05	2.77E-07	1.57E-08	6.63	61.9	3.36E-09	1.76E-10	1.29E-07	8.05E-09	38421.3	3135.9
	2	0.05	2	0.05	1.29E-07	7.06E-09	6.55	60.5	1.88E-09	1.34E-10	6.36E-08	3.89E-09	33889.3	3189.5
	1	0.05	2	0.05	4.61E-08	2.70E-09	5.39	60.4	6.43E-10	1.05E-10	2.27E-08	1.47E-09	35382.4	6219
	9-BNF	/			2.21E-05	1.30E-06	4.68	61	2.29E-05	1.05E-06	/	/	/	/
	9-BF	/			2.21E-05	1.30E-06	4.68	61	2.28E-05	1.09E-06	/	/	/	/
	8-BNF	/			4.61E-08	2.64E-09	5.67	61.5	1.06E-08	6.30E-10	/	/	/	/
	8-BF	/			4.61E-08	2.64E-09	5.67	61.5	/	/	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 30. Experimental results of the sorption isotherms of U(VI) onto illite at 60 ºC in atmospheric pCO2 in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.05	2	0.05	1.73E-05	9.08E-07	7.41	57.7	7.06E-07	8.40E-08	8.29E-06	2.17E-07	11738.1	1430.9
	6	0.049	1.96	0.05	6.48E-06	3.36E-07	7.57	57	1.16E-07	3.87E-08	3.25E-06	8.68E-08	28098.3	9430.2
	5	0.049	1.96	0.05	2.16E-06	1.08E-07	7.65	57.7	5.11E-07	2.98E-08	8.41E-07	2.67E-08	1644.6	109.2
	4	0.052	2.08	0.05	3.17E-07	1.74E-08
					2.64E-08	1.45E-09	7.49	56.7	2.20E-08	1.09E-09	/	/	/	/
	8-BF	/	/	/	2.64E-08	1.45E-09	7.49	56.7	2.17E-08	1.13E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 31. Experimental results of the sorption isotherms of U(VI) onto illite at 60 ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.053	2.12	0.05	1.73E-05	9.08E-07	7.4	57.6	4.99E-06	2.44E-07	5.80E-06	4.66E-07	1161	109.2
	6	0.051	2.04	0.05	6.48E-06	3.36E-07	7.42	57.7	4.35E-06	2.18E-07	1.04E-06	1.98E-07	240.1	47.2
	5	0.05	2	0.05	2.16E-06	1.08E-07	7.45	57.3	1.21E-06	5.88E-08	4.73E-07	6.27E-08	389.3	54.9
	4	0.051	2.04	0.05	6.42E-07	3.44E-08	7.45	57.8	2.26E-07	1.18E-08	2.04E-07	1.86E-08	899.1	94.3
	3	0.05	2	0.05	3.21E-07	1.70E-08	7.45	59	8.15E-08	5.04E-09	1.20E-07	9.38E-09	1468.6	146.6
	2	0.051	2.04	0.05	1.50E-07	7.62E-09	7.48	59.3	2.47E-08	1.55E-09	6.13E-08	4.12E-09	2481.3	228.3
	1	0.05	2	0.05	5.35E-08	2.95E-09	7.47	58.2	6.34E-09	5.46E-10	2.36E-08	1.62E-09	3715.2	408.9
	9-BNF	/	/	/	1.73E-05	9.52E-07	7.39	55.8	1.78E-05	8.82E-07	/	/	/	/
	9-BF	/	/	/	1.73E-05	9.52E-07	7.39	55.8	1.80E-05	8.40E-07	/	/	/	/
	8-BNF	/	/	/	5.35E-08	2.87E-09	7.48	58.7	5.95E-08	3.40E-09	/	/	/	/
	8-BF	/	/	/	5.35E-08	2.87E-09	7.48	58.7	5.92E-08	2.86E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 32. Experimental results of the sorption isotherms of U(VI) onto illite at 60 ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.049	1.96	0.05	2.36E-05	1.40E-06	6.82	59.7	1.97E-06	1.51E-07	1.10E-05	7.72E-07	581.7	3.7
	6	0.051	2.04	0.05	8.85E-06	5.19E-07	6.83	59.9	8.19E-07	6.30E-08	3.94E-06	2.75E-07	499.2	3.7
	5	0.049	1.96	0.05	2.95E-06	1.68E-07	6.83	59.8	2.72E-07	2.27E-08	1.37E-06	9.35E-08	541.1	3.7
	4	0.049	1.96	0.05	3.19E-07	1.65E-08
					2.66E-08	1.38E-09	6.85	59.4	2.48E-08	1.26E-09	/	/	/	/
	8-BF	/	/	/	2.66E-08	1.38E-09	6.85	59.4	2.45E-08	1.30E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 33. Experimental results of the sorption isotherms of U(VI) onto illite at 80 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.049	1.96	0.05	2.21E-05	1.20E-06	6.33	79.7	3.99E-08	1.68E-09	1.13E-05	6.78E-07	282196.6	20739.2
	6	0.051	2.04	0.05	8.29E-06	4.45E-07	6.59	78.8	2.99E-08	1.34E-09	4.05E-06	2.41E-07	135606.4	10115.2
	5	0.05	2	0.05	2.76E-06	1.43E-07	6.71	78.8	1.90E-08	8.82E-10	1.37E-06	7.98E-08	72287.6	5379
	4	0.05	2	0.05	5.54E-07	3.17E-08	6.58	79.3	5.51E-09	2.77E-10	2.74E-07	1.73E-08	49713.4	4016.6
	3	0.05	2	0.05	2.77E-07	1.57E-08	6.5	78.7	1.73E-09	1.39E-10	1.38E-07	8.59E-09	79451.5	8071
	2	0.05	2	0.05	1.29E-07	7.06E-09	6.61	80.3	4.69E-10	1.93E-11	6.44E-08	3.90E-09	137118.6	10039.4
	1	0.05	2	0.05	4.61E-08	2.70E-09	6.49	80	1.48E-10	6.72E-12	2.30E-08	1.47E-09	155024.1	12174.6
	9-BNF	/	/	/	2.21E-05	1.30E-06	4.56	78.6	2.04E-05	8.82E-07	/	/	/	/
	9-BF	/	/	/	2.21E-05	1.30E-06	4.58	78.6	2.08E-05	9.24E-07	/	/	/	/
	8-BNF	/	/	/	4.61E-08	2.64E-09	5.77	79.2	2.00E-09	4.08E-10	/	/	/	/
	8-BF	/	/	/	4.61E-08	2.64E-09	5.77	79.2	1.63E-09	3.95E-10	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Illite	S/L		UVI) ini.		pH	T	eq. U(VI)		sorb. U(VI)		KD

Table C . 34. Experimental results of the sorption isotherms of U(VI) onto illite at 80 ºC in atmospheric pCO2 in 0.1 M of NaCl at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.05	2	0.05	1.73E-05	9.08E-07	7.46	76.6	6.74E-08	3.15E-09	2.04E-07	1.36E-08	3026.5	246.7
	6	0.049	1.96	0.05	6.48E-06	3.36E-07	7.45	71.9	5.11E-08	2.65E-09	6.50E-08	5.19E-09	1272.4	121.2
	5	0.049	1.96	0.05	2.16E-06	1.08E-07	7.54	76.5	/	/	/	/	/	/
	4	0.052	2.08	0.05	3.17E-07	1.74E-08	7.63	72.9	5.17E-09	5.04E-10	5.50E-08	3.45E-09	10635.6	1233.9
	3	0.05	2	0.05	1.58E-07	8.60E-09	7.55	73.4	1.47E-09	4.20E-10	2.91E-08	1.80E-09	19814.3	5791.7
	2	0.05	2	0.05	7.39E-08	3.86E-09	7.63	73.1	5.59E-10	1.76E-10	1.37E-08	8.13E-10	24447.4	7856
	1	0.049	1.96	0.05	2.64E-08	1.49E-09	7.58	76.5	2.44E-10	1.72E-10	4.96E-09	3.26E-10	20337.8	14438.8
	9-BNF	/	/	/	1.73E-05	9.74E-07	6.76	20.3	5.48E-07	2.82E-08	/	/	/	/
	9-BF	/	/	/	1.73E-05	9.74E-07	6.76	20.3	4.76E-07	2.39E-08	/	/	/	/
	8-BNF	/	/	/	2.64E-08	1.45E-09	7.3	20.6	1.01E-08	6.30E-10	/	/	/	/
	8-BF	/	/	/	2.64E-08	1.45E-09	7.3	20.6	9.96E-09	8.82E-10	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 35. Experimental results of the sorption isotherms of U(VI) onto illite at 80 ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	7	0.052	2.08	0.05	1.73E-05	9.08E-07	7.63	73.6	2.51E-06	1.47E-07	7.10E-06	4.76E-07	2830.7	252.2
	6	0.053	2.12	0.05	6.48E-06	3.36E-07	7.59	78.3	3.77E-06	2.14E-07	1.28E-06	1.91E-07	338.3	54.1
	5	0.049	1.96	0.05	2.16E-06	1.08E-07	7.6	75.2	1.05E-06	5.04E-08	5.68E-07	6.26E-08	543	65.3
	4	0.048	1.92	0.05	6.42E-07	3.44E-08	7.67	78.7	2.32E-07	1.09E-08	1.36E-07	1.53E-08	586.2	71.6
	3	0.051	2.04	0.05	3.21E-07	1.70E-08	7.66	79.5	7.77E-08	4.62E-09	8.25E-08	7.10E-09	1061.5	111
	2	0.05	2	0.05	1.50E-07	7.62E-09	7.66	79.1	2.46E-08	1.39E-09	4.51E-08	3.22E-09	1831.7	166.5
	1	0.052	2.08	0.05	5.35E-08	2.95E-09	7.64	77.7	5.04E-09	5.04E-10	1.73E-08	1.19E-09	3429.5	416.8
	9-BNF	/	/	/	1.73E-05	9.52E-07	7.64	73.6	1.17E-05	5.46E-07	/	/	/	/
	9-BF	/	/	/	1.73E-05	9.52E-07	7.64	73.6	1.16E-05	5.46E-07	/	/	/	/
	8-BNF	/	/	/	5.35E-08	2.87E-09	7.7	77	4.20E-08	2.06E-09	/	/	/	/
	8-BF	/	/	/	5.35E-08	2.87E-09	7.7	77	4.10E-08	2.18E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 36. Experimental results of the sorption isotherms of U(VI) onto illite at 80 ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	7	0.05	2	0.05	2.36E-05	1.40E-06	7.02	79.7	3.04E-06	1.68E-07	1.03E-05	7.52E-07	3383.8	310.3
	6	0.05	2	0.05	8.85E-06	5.19E-07	7.03	79.8	1.24E-06	6.30E-08	3.80E-06	2.79E-07	3069.5	273.9
	5	0.05	2	0.05	2.95E-06	1.68E-07	7.02	79.3	8.91E-07	4.62E-08	1.03E-06	9.11E-08	1155.7	118.6
	4	0.05	2	0.05	3.19E-07	1.65E-08	7.06	79.5	3.63E-08	2.14E-09	5.27E-08	4.06E-09	1451.4	140.8
	3	0.049	1.96	0.05	1.59E-07	8.15E-09	7.05	79.3	1.53E-08	8.40E-10	2.84E-08	2.04E-09	1859.5	168.3
	2	0.052	2.08	0.05	7.44E-08	3.64E-09	7.05	80.1	4.73E-09	2.69E-10	1.36E-08	8.59E-10	2881.6	244.8
	1	0.05	2	0.05	2.66E-08	1.42E-09	7.05	77.6	1.22E-09	1.30E-10	5.29E-09	3.49E-10	4344.8	545.7
	9-BNF	/	/	/	2.36E-05	1.26E-06	7.01	77.8	1.16E-05	5.88E-07	/	/	/	/
	9-BF	/	/	/	2.36E-05	1.26E-06	7.01	77.8	1.13E-05	5.46E-07	/	/	/	/
	8-BNF	/	/	/	2.66E-08	1.38E-09	7.06	77.4	1.23E-08	7.14E-10	/	/	/	/
	8-BF	/	/	/	2.66E-08	1.38E-09	7.06	77.4	1.18E-08	5.88E-10	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Illite	S/L		U(VI) ini.		pH	T	U(VI) eq.		U(VI) sorb.		KD

Table C . 37. Experimental results of the sorption isotherms of U(VI) onto illite at 80 ºC in 5% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	C.5.2. Sorption edges:
	7	0.051	2.04	0.05	2.16E-05	1.20E-06	6.84	81.8	1.88E-05	9.24E-07	1.38E-06	7.42E-07	73.53	39.56
	6	0.05	2	0.05	8.12E-06	4.43E-07	6.85	81.3	7.05E-06	3.49E-07	5.36E-07	2.82E-07	76.03	40.24
	5	0.05	2	0.05	2.71E-06	1.43E-07	6.86	80.2	2.16E-06	1.18E-07	2.71E-07	9.28E-08	125.24	43.43
	4	0.049	1.96	0.05	3.19E-07	1.65E-08	6.86	76.7	2.47E-07	1.26E-08	3.65E-08	1.06E-08	147.86	43.73
	3	0.051	2.04	0.05	1.59E-07	8.15E-09	6.84	81.2	1.16E-07	6.72E-09	2.13E-08	5.21E-09	183.29	46.14
	2	0.05	2	0.05	7.44E-08	3.64E-09	6.84	81.6	5.02E-08	3.36E-09	1.21E-08	2.50E-09	241.04	52.32
	1	0.051	2.04	0.05	2.66E-08	1.42E-09	6.84	81.3	1.82E-08	1.68E-09	4.08E-09	1.08E-09	223.64	62.83
	9-BNF	/	/	/	2.16E-05	1.15E-06	6.8	80	2.12E-05	1.01E-06	/	/	/	/
	9-BF	/	/	/	2.16E-05	1.15E-06	6.8	80	2.13E-05	1.01E-06	/	/	/	/
	8-BNF	/	/	/	2.66E-08	1.38E-09	6.79	79.1	2.45E-08	1.81E-09	/	/	/	/
	8-BF	/	/	/	2.66E-08	1.38E-09	6.79	79.1	2.49E-08	1.85E-09	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
					ini.				eq.		sorb.
	ID	Illite	S/L		U(VI)		pH	T	U(VI)		U(VI)		KD

Table C . 38. pH results of each tube measured before the batch equilibrium at 20 ºC in atmospheric pCO2.

 C

	Tubes	S/L	pH of NaCl equilibrated HCl/NaOH (0.01M) pH final
		g/L		µL
	edge 1	1	3.89	-	3.9
	edge 2	1	5.6+HCl	60µL	4.5
	edge3	0.5	5.6	-	5.6
	edge4	0.5	7+HCl	60µL	6.5
	edge5	0.5	7	-	7
	edge6	0.5	7.7	-	7.7
	edge7	1	7.7+NaOH	60µL	8.5
	edge8	1	9.8	-	9.8
	Blanc 9	0	3.89	-
	Blanc 10	0	7	-
	Blanc 11	0	9.8	-

Table C . 39. pH results of each tube measured before the batch equilibrium at 80 ºC in atmospheric pCO2.

 C

	ID	Illite	S/L		pH	Temp	U(VI) eq.		U(VI) sorb.	KD
		(g)	g/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	10-BF	/	/	/ 9.16	76.2 3.66E-08 1.76E-09
	10-BNF	/	/	/ 9.16	76.2 3.50E-08 1.55E-09
	1	0.025	1 0.04 5.01	79.7 1.17E-09 1.13E-10 2.95E-08 1.99E-09	25150.5	2966.7
	2	0.025	1 0.04 6.54	80.7 1.18E-10 2.35E-11 3.05E-08 2.01E-09 257723.4	53929.7
	3	0.013 0.52 0.04	6.7	80.5 5.46E-10 9.24E-11 5.79E-08 5.42E-09 106005.9	20497.6
	4	0.014 0.56 0.04 7.19	79.2 3.91E-13 5.88E-14 5.47E-08 4.85E-09	1.4E+08 24464341
	5	0.015	0.6 0.04	7.4	77.5 3.40E-13 5.46E-14 5.11E-08 4.33E-09	1.5E+08 27244360
	6	0.013 0.52 0.04 8.19	76.3 6.37E-12 3.32E-13 5.89E-08 5.48E-09	9247101	985716.4
	7	0.013 0.52 0.04	8.2	78.7 7.34E-12 4.03E-13 5.89E-08 5.48E-09	8024833	866807.8
	C.6
	Tubes	S/L		pH of NaCl equilibrated HCl/NaOH (0.01M) pH final
		g/L					µL
	edge 1	1		/			/		3.89
	edge 2	1		HCl 0,01 mol/L	0.0600	6.05
	edge3	0.5		/			/		6.05
	edge4	0.5		HCl 0,01 mol/L	0.0600	7.22
	edge5	0.5		/			/		7.22
	edge6	0.5		/			/		7.99
	edge7	1		NaOH 0,01 mol/L	0.0600	7.99
	edge8	1		/			/		9.84
	Blanc 9	0		/			/		3.89
	Blanc 10	0		/			/		7.22
	Blanc 11	0		/			/		9.84

. Callovo-Oxfordian clay fraction: sorption isotherm experimentsTable C . 44. Alkalinity measurements in sorption isotherms experiments of U(VI)/clay fraction system at different temperatures. *synthetic Callovo-Oxfordian porewater.

 C

	atmosphere	T	Illite	pH	Acid H2SO4	Vol. acid	±	[CO3 2-/HCO3 -]	±
		ºC	g		N	digits		meq/L
	Ar(g)	20.5	0.049	7.06	0.1	89	35.6	2.23E-04	8.90E-05
	1% CO2(g)	21.5	0.059	7.24	0.16	118	47.2	2.36E-03	9.44E-04
	1% CO2(g)*	21.5	0.05	7.24	0.16	118	47.2	2.36E-03	9.44E-04
	Ar(g)	77.5	0.05	7.08	0.16	89	35.6	2.23E-04	8.90E-05
	5% CO2(g)	67.	0.05	7.15	0.16	115	46	2.30E-03	9.20E-04
	5% CO2(g)*	68	0.049	7.08	0.16	168	67.2	3.36E-03	1.34E-03

Table C . 45. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio.

 C

	2	0.05	2	0.05	1.29E-07	7.06E-09	7.08	21.4	7.02E-09	4.62E-10	6.11E-08	3.87E-09	8704.8	795.4
	1	0.052	2.08	0.05	4.61E-08	2.70E-09	7.09	21.4	2.60E-09	1.55E-10	2.09E-08	1.40E-09	8047.3	722.9
	9-BNF	/	/	/	4.61E-08	2.70E-09	5.46	21	3.54E-08	1.68E-09	/	/	/	/
	9-BF	/	/	/	4.61E-08	2.70E-09	5.46	21	4.25E-08	1.93E-09	/	/	/	/
	8-BNF	/	/	/	2.21E-05	1.20E-06	4.75	21	2.22E-05	1.01E-06	/	/	/	/
	8-BF	/	/	/	2.21E-05	1.20E-06	4.75	21	2.23E-05	1.09E-06	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		ini U(VI)		pH	T	eq. U(VI)		sorb. U(VI)		KD
	7	0.051	2.04	0.05	2.21E-05	1.20E-06	6.69	21.2	1.30E-06	9.66E-08	1.02E-05	6.44E-07	7860.6	767.2
	6	0.049	1.96	0.05	8.29E-06	4.45E-07	6.98	21.2	3.41E-07	1.68E-08	4.06E-06	2.50E-07	11908.3	940.6
	5	0.05	2	0.05	2.76E-06	1.43E-07	7.13	21.2	7.90E-08	5.04E-09	1.34E-06	7.95E-08	17000	1480
	4	0.05	2	0.05	5.54E-07	3.17E-08	7.11	21.2	2.63E-08	1.55E-09	2.64E-07	1.72E-08	10040.8	885.6
	3	0.049	1.96	0.05	2.77E-07	1.57E-08	7.08	21.3	1.78E-08	1.01E-09	1.32E-07	8.72E-09	7436	646.9

Table C . 46. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20 ºC in 1% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	9-BF	/	/	/	5.35E-08	2.95E-09	7.22	20.8	6.75E-08	2.98E-09	/	/	/	/
	8-BF	/	/	/	1.73E-05	9.08E-07	7.17	21	1.87E-05	1.05E-06	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		ini. U(VI)		pH	T	eq. U(VI)		sorb. U(VI)		KD
	7	0.05048	2.02	0.05	1.17E-05	6.13E-07	7.21	21	6.95E-06	4.12E-07	2.82E-06	3.90E-07	405.6	61
	6	0.05167	2.07	0.05	6.48E-06	3.36E-07	7.22	20.9	3.67E-06	1.85E-07	1.62E-06	2.00E-07	442	58.7
	5	0.05085	2.03	0.05	2.16E-06	1.10E-07	7.23	20.9	9.50E-07	4.62E-08	6.85E-07	6.41E-08	721.1	76.1
	4	0.04943	1.98	0.05	6.42E-07	3.44E-08	7.23	20.9	2.42E-07	1.34E-08	2.87E-07	2.35E-08	1189.7	117.5
	3	0.05085	2.03	0.05	3.21E-07	1.70E-08	7.23	20.7	9.45E-08	5.88E-09	1.53E-07	1.12E-08	1614.3	155.4
	2	0.05004	2	0.05	1.50E-07	7.62E-09	7.23	20.8	3.91E-08	2.44E-09	7.49E-08	5.10E-09	1916.6	177
	1	0.0487	1.95	0.05	5.35E-08	2.95E-09	7.23	20.7	1.02E-08	9.66E-10	2.94E-08	2.03E-09	2880	337.3

Table C . 47. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 20 ºC in 1% of CO2(g) in synthetic Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	8-BF	/	/	/	1.73E-05	9.08E-07	7.17	21	1.87E-05	1.05E-06	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		ini. U(VI)		pH	T	eq. U(VI)		sorb. U(VI)		KD
	7	0.05109	2.044	0.05	1.17E-05	6.13E-07	7.08	21.2	1.16E-05	5.46E-07	4.94E-07	4.21E-07	42.5	36.2
	6	0.05053	2.021	0.05	6.48E-06	3.36E-07	7.1	21.1	6.46E-06	3.40E-07	2.82E-07	2.47E-07	43.6	38.3
	5	0.05065	2.026	0.05	2.16E-06	1.10E-07	7.12	21.2	2.24E-06	1.09E-07	5.29E-08	8.00E-08	23.7	35.8
	4	0.05006	2.002	0.05	6.42E-07	3.44E-08	7.12	21.3	6.44E-07	2.90E-08	8.30E-08	2.61E-08	129	41
	3	0.04996	1.998	0.05	3.21E-07	1.70E-08	7.13	21.3	3.16E-07	1.60E-08	4.45E-08	1.34E-08	140.9	43
	2	0.04989	1.996	0.05	1.50E-07	7.62E-09	7.13	21.2	1.47E-07	6.72E-09	2.10E-08	5.89E-09	142.8	40.6
	1	0.05002	2	0.05	5.35E-08	2.95E-09	7.11	21.1	6.48E-08	2.98E-09	1.32E-09	2.38E-09	20.4	36.8
	9-BF	/	/	/	5.35E-08	2.95E-09	7.22	20.8	6.75E-08	2.98E-09	/	/	/	/

Table C . 48. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in Ar(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio of 2 g/L. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

					2.21E-05	1.20E-06	4.06	77.6	1.84E-05	8.82E-07	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		ini. U(VI)		pH	T	eq. U(VI)		sorb. U(VI)		KD
	7	0.05	2	0.05	2.21E-05	1.20E-06	7.08	77.1	1.22E-06	5.46E-08	1.04E-05	6.58E-07	8575.9	662.7
	6	0.05	2	0.05	8.29E-06	4.45E-07	7.17	77.9	4.59E-07	2.14E-08	3.92E-06	2.44E-07	8530.2	664.3
	5	0.053	2.12	0.05	2.76E-06	1.43E-07	7.2	78.5	1.38E-07	9.66E-09	1.24E-06	7.43E-08	8962.3	825.7
	4	0.05	2	0.05	5.54E-07	3.00E-08	7.25	77.9	2.88E-08	1.68E-09	2.62E-07	1.65E-08	9117.5	781.8
	3	0.052	2.08	0.05	2.77E-07	1.48E-08	7.25	78.8	1.31E-08	7.98E-10	1.27E-07	7.82E-09	9670.9	838.3
	2	0.05	2	0.05	1.29E-07	6.66E-09	7.23	78.5	1.04E-08	7.56E-10	5.94E-08	3.68E-09	5723.5	547.4
	1	0.05	2	0.05	4.61E-08	2.57E-09	7.14	78.3	2.34E-09	1.43E-10	2.19E-08	1.40E-09	9338.7	826.1
	9-BNF	/	/	/	4.61E-08	2.57E-09	5.82	78.6	8.99E-10	3.70E-10	/	/	/	/
	9-BF	/	/	/	4.61E-08	2.57E-09	5.82	78.6	4.54E-10	3.61E-10	/	/	/	/

Table C . 49. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in 5% of CO2(g) in 0.1 M of NaNO3 at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C

	9-BF	/	/	/	9.96E-09	6.10E-10	7.1	76.8	5.16E-09	3.78E-10	/	/	/	/
	8-BNF	/	/	/	1.57E-05	8.31E-07	7.18	75.7	1.53E-05	7.14E-07	/	/	/	/
	8-BF	/	/	/	1.57E-05	8.31E-07	7.18	75.7	1.48E-05	7.14E-07	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		U(VI) ini.		pH	T	U(VI) eq.		U(VI) sorb.		KD
	7	0.049	1.96	0.05	1.06E-05	5.61E-07	7.21	74.5	3.87E-06	1.89E-07	3.44E-06	3.15E-07	889.8	92.3
	6	0.049	1.96	0.05	5.90E-06	3.08E-07	7.22	75.3	1.86E-06	1.13E-07	2.06E-06	1.76E-07	1110.4	116.4
	5	0.05	2	0.05	1.97E-06	9.89E-08	7.19	76.7	3.47E-07	2.44E-08	8.10E-07	5.50E-08	2336.4	228.4
	4	0.051	2.04	0.05	1.19E-07	7.16E-09	7.23	76	8.03E-09	1.09E-09	5.46E-08	3.81E-09	6808.8	1041.3
	3	0.05	2	0.05	5.97E-08	3.55E-09	7.21	77.7	2.70E-09	3.07E-10	2.85E-08	1.92E-09	10557.5	1394.1
	2	0.052	2.08	0.05	2.79E-08	1.60E-09	7.22	78.5	1.82E-09	3.74E-10	1.25E-08	8.51E-10	6904.4	1497.7
	1	0.052	2.08	0.05	9.96E-09	6.10E-10	7.25	78.6	6.43E-10	1.26E-10	4.48E-09	3.20E-10	6966.4	1453.6
	9-BNF	/	/	/	9.96E-09	6.10E-10	7.1	76.8	5.50E-09	4.62E-10	/	/	/	/

Table C . 50. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian clay fraction at 80 ºC in 5% of CO2(g) in Callovo-Oxfordian porewater at fixed solid-liquid ratio. (BF: blank sample filtrated; BNF: blank sample no filtrated)

 C Impact of Increasing the Temperature up to 80 ºC on the Behaviour of Radionuclides in Callovo-Oxfordian Formation: Application to Uranium

	9-BF				9.96E-09	6.10E-10	6.9	73.6	8.70E-09	6.30E-10	/	/	/	/
	8-BNF				1.57E-05	8.31E-07	6.93	74.3	1.55E-05	6.72E-07	/	/	/	/
	8-BF				1.57E-05	8.31E-07	6.93	74.3	1.53E-05	6.72E-07	/	/	/	/
		(g)	g/L	±	mol/L	±	/	ºC	mol/L	±	mol/g	±	mL/g	±
	ID	Clay fraction	S/L		U(VI) ini.		pH	T	U(VI) eq.		U(VI) sorb.
	7	0.052	2.08	0.05	1.06E-05	5.61E-07	7.1	74	9.71E-06	4.62E-07	4.39E-07	3.50E-07	45.2	36.1
	6	0.049	1.96	0.05	5.90E-06	3.08E-07	7.07	76.5	5.08E-06	2.31E-07	4.18E-07	1.97E-07	82.3	38.9
	5	0.051	2.04	0.05	1.97E-06	9.89E-08	7.07	77.4	1.72E-06	7.98E-08	1.19E-07	6.24E-08	69.3	36.4
	4	0.051	2.04	0.05	1.19E-07	7.16E-09	7.09	76.4	8.82E-08	5.46E-09	1.53E-08	4.43E-09	173.7	51.4
	3	0.049	1.96	0.05	5.97E-08	3.55E-09	7.16	77.2	4.35E-08	2.73E-09	8.30E-09	2.29E-09	190.8	54.1
	2	0.05	2	0.05	2.79E-08	1.60E-09	7.17	77.2	1.69E-08	1.51E-09	5.50E-09	1.11E-09	325.4	71.9
	1	0.05	2	0.05	9.96E-09	6.10E-10	7.16	77.6	6.97E-09	4.20E-10	1.49E-09	3.72E-10	213.9	54.9
	9-BNF				9.96E-09	6.10E-10	6.9	73.6	8.74E-09	6.72E-10	/	/	/	/

KD

C.7. Callovo-Oxfordian claystone samples: isotherm experiments

Table C . 51. Alkalinity measurements in sorption isotherms experiments of U(VI)/claystone system at 20 and 80ºC in synthetic Callovo-Oxfordian porewater.

 C

	atmosphere T	pH Acid H2SO4 Vol. acid	±	Alkalinity	±
		ºC	N		digits	meq/L
	1% CO2(g)	20 7.42	0.16	164 65.6	3.28 1.31
	5% CO2(g)	80 6.85	0.16	108 43.2	2.16 0.84

Table C . 52. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian claystone at 20 ºC in 1% of CO2(g) in Cox porewater at fixed solid-liquid ratio.

 C

					2.40E-05		1.24E-06	7.27	18.9	2.40E-05		1.09E-06	/	/	/	/
		(g)	g/L	±	mol/	L	±	/	ºC	mol/	L	±	mol/g	±	mL/g	±
					initial					eq.			sorb.
	ID	COx	S/L		U(VI)			pH	Temp	U(VI)			U(VI)		KD
	8	0.2	39.4	0.6	9.61E-06		4.95E-07	7.51	18.2	8.41E-06		3.78E-07	3.06E-08	9.61E-09	3.64	1.15
	7	0.075	39.4	0.6	3.61E-06		1.83E-07	7.5	18.2	3.20E-06		1.51E-07	1.04E-08	3.85E-09	3.24	1.21
	6	0.03	39.5	0.6	1.44E-06		6.94E-08	7.5	18.4	1.24E-06		5.88E-08	5.02E-09	1.49E-09	4.03	1.21
	5	0.4	39.4	0.6	9.02E-07		4.58E-08	7.5	18.1	7.83E-07		3.53E-08	3.02E-09	8.96E-10	3.85	1.16
	4	0.2	39.4	0.6	4.51E-07		2.26E-08	7.5	18.1	3.87E-07		1.72E-08	1.64E-09	4.38E-10	4.24	1.15

Table C . 53. Experimental results of the sorption isotherms of U(VI) onto Callovo-Oxfordian claystone samples at 80 ºC in 5% of CO2(g) in Callovo-Oxfordian porewater at fixed solid-liquid ratio.

 C

	7	0.051	2.04	0.05	2.16E-05	1.20E-06	6.84	81.8	1.88E-05	9.24E-07	1.38E-06	7.42E-07	73.5	39.6
	6	0.05	2	0.05	8.12E-06	4.43E-07	6.85	81.3	7.05E-06	3.49E-07	5.36E-07	2.82E-07	76	40.2
	5	0.05	2	0.05	2.71E-06	1.43E-07	6.86	80.2	2.16E-06	1.18E-07	2.71E-07	9.28E-08	125.2	43.4
	4	0.049	1.96	0.05	3.19E-07	1.65E-08	6.86	76.7	2.47E-07	1.26E-08	3.65E-08	1.06E-08	147.9	43.7
	3	0.051	2.04	0.05	1.59E-07	8.15E-09	6.84	81.2	1.16E-07	6.72E-09	2.13E-08	5.21E-09	183.3	46.1
	2	0.05	2	0.05	7.44E-08	3.64E-09	6.84	81.6	5.02E-08	3.36E-09	1.21E-08	2.50E-09	241	52.3
	1	0.051	2.04	0.05	2.66E-08	1.42E-09	6.84	81.3	1.82E-08	1.68E-09	4.08E-09	1.08E-09	223.6	62.8

Table D . 2. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 20 ºC and 1% of CO2(g)

 D

		Experimental		Modelled
	S/L		U(VI) sol
	mL/g	±	µmol/L	±	µM
	1.2476	0.06238	0.087395 0.007983	0.093707
	2.6596	0.13298	0.04916	0.007983	0.042252
	5.072667	0.370841	0.026471 0.008824	0.022362
	10.61333	0.775895	0.013866 0.007143	0.011844
	15.15733	1.108088	0.011345 0.007143	0.008081
	20.939	2.186096	0.009244 0.007143	0.006135
	25.824	2.696105	0.006429 0.002731	0.005009

Table D . 3. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 40 ºC and 1% of CO2(g)

 D

		Experimental		Modelled
	S/L		U(VI) sol
	mL/g	±	µmol/L	±	µM
	1.0308	0.05154	0.115126 0.008403	0.035137
	2.5324	0.12662	0.059664 0.010504	0.092077
	5.068	0.3705	0.034034 0.006303	0.225963
	10.038	0.733835	0.017227 0.005882	0.45691
	25.31818	2.667801	0.012185 0.006303	1.518603

Table D . 4. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 60 ºC and 3% of CO2(g)

 D

		Experimental		Modelled
	S/L		U(VI) sol		U(VI) sol
	mL/g	±	µmol/L	±	µM
	1.0244	0.05122	0.110084 0.006723	0.031181
	2.5008	0.12504	0.062941 0.003992	0.07871
	4.958667	0.362507	0.035588 0.003067	0.199101
	9.824667	0.718239	0.024916 0.002941	0.400656
	25.51	2.663322	0.017059 0.002815	1.37519

Table D . 5. Modelling results using PEST coupled to Phreeqc for experiments as a function of S/L in the absence of Ca at 80 ºC and 5% of CO2(g)

 D

	Experimental				Modelled
	S/L		U(VI) sol		U(VI) sol
	mL/g	±	µmol/L	±	µM
	1.016	0.0508	0.147899 0.007563	0.017524
	2.496644	0.123794	0.090756 0.005042	0.049778
	5.074	0.370938	0.048109 0.003319	0.164974
	9.934667	0.726281	0.055	0.003487	0.318007
	14.786	1.080941	0.055798 0.003319	0.472435
	19.618	2.048179	0.045168 0.002983	0.916948
	24.639	2.572387	0.035294 0.002983

D.2.2. Type II experiments (ΔCa):

Table D . 6. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 20 ºC and 1% of CO2(g)

 D

		Experimental	Modelled
	Ca	U(VI) sol	U(VI) sol
	mM	µM	±	µM
	0	0.044538	0.016807	0.016043
	0.1	0.081513	0.015966	0.015756
	0.2	0.094118	0.015966	0.015512
	0.3	0.107143	0.015546	0.015625
	0.4	0.128151	0.015546	0.015514
	0.5	0.133613	0.015966	0.015399
	1	0.193697	0.018067	0.015589
	2	0.24958	0.018067	0.015566
	3	0.316807	0.02605	0.017134
	4	0.369328	0.036134	0.019805
	5	0.401681	0.037815	0.020129

Table D . 7. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 40 ºC and 1% of CO2(g)

 D

		Experimental	Modelled
	Ca	U(VI) sol	U(VI) sol
	mM	µM	±	µM
	0	0.051681	0.011345	0.030333
	0.1	0.060924	0.013025	0.031205
	0.2	0.07605	0.014286	0.03097
	0.3	0.086975	0.015126	0.031266
	0.4	0.097899	0.015966	0.030772
	0.5	0.108403	0.017227	0.031409
	1	0.159244	0.020168	0.03137
	2	0.244538	0.033613	0.033017
	3	0.294118	0.042017	0.034205
	4	0.32605	0.039916	0.033902
	5	0.361345	0.046218	0.035633

Table D . 8. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 60 ºC and 3% of CO2(g)

 D

	0.2	0.047479	0.007983	0.01524
	0.3	0.055882	0.007563	0.015208
	0.4	0.063866	0.009664	0.014757
	0.5	0.070588	0.009664	0.015334
	1	0.101261	0.012185	0.015385
	2	0.159244	0.015546	0.015763
	3	0.217647	0.015126	0.01539
	4	0.266387	0.015966	0.015445
	5	0.310084	0.019748	0.016072
		Experimental	Modelled
	Ca	U(VI) sol	U(VI) sol
	mM	µM	±	µM
	0	0.040756	0.008403	0.014631
	0.1	0.055462	0.008824	0.015316

Table D . 9. Modelling results using PEST coupled to Phreeqc for experiments as a function of Ca at fixed S/L, 80 ºC and 5% of CO2(g)

 D

	SOLUTION 5
		temp		20
		pH		7.2
		units		mol/kgw
		Alkalinity 2.25e-3
		U		5.21e-7
		Na 0.1
		N(5) 0.1 charge
		-water	1 # kg
	EXCHANGE 5
	XNO3 3.93e-2
	EQUILIBRIUM_PHASES 5
	fix_pH -7.2 NaOH 10
	Clarkeite 0 0
	END
	SOLUTION 6
		temp		20
		pH		7.18
		units		mol/kgw
		Alkalinity 2.25e-3
		U		5.21e-7
		Na 0.1
		N(5) 0.1 charge
		-water	1 # kg
	EXCHANGE 6
	XNO3 5.24e-2
	EQUILIBRIUM_PHASES 5
	fix_pH -7.18 NaOH 10
	Clarkeite 0 0
	END
	SOLUTION 7
		temp pH		Experimental 20 7.14	Modelled
	Ca	units		U(VI) sol mol/kgw	U(VI) sol
	mM Alkalinity 2.25e-3 µM	±	µM
	0	U	0.062185 5.21e-7	0.004622	0.015296
	0.1 0.2 0.3 EXCHANGE 7 0.068487 Na 0.1 0.071849 N(5) 0.1 charge 0.07395 -water 1 # kg	0.005042 0.004622 0.005042	0.015409 0.01536 0.015075
	0.4 XNO3 6.55e-2 0.078992	0.005042	0.015228
	0.5 EQUILIBRIUM_PHASES 7 0.083613	0.005042	0.01534
	1 fix_pH -7.14 NaOH 10 0.109244	0.005882	0.015218
	2 Clarkeite 0 0 0.172269 3 0.22437 END	0.008403 0.010084	0.014924 0.015134
	4		0.286975	0.015126	0.01568
	5 Control File:	0.320168	0.015966	0.015595

80 ºC and 5% of CO2(g) [U(VI)] as a function of calcium

	* observation groups
	U_data
	* observation data
	U_1	8.73950E-08 1	U_data
	U_2	4.91597E-08 1	U_data
	U_3	2.64706E-08 1	U_data
	U_4	1.38655E-08 1	U_data
	U_5	1.13445E-08 1	U_data
	U_6	9.24370E-09 1	U_data
	U_7	6.42857E-09 1	U_data
	* model command line
	phreeqc fittingE1A.pqi fittingE1A.pqo ThermoChimie_PHREEQC_Davies_v9b0.dat
	* model input/output
	fittingE1A.tpl fittingE1A.pqi
	fittingE1A.ins fittingE1A.prn
	Template file:

  total num sites or density (mol/m2); surface are (m2/g); g of mineral (considering 100 g of rock and all mineral is reactive) Surfill_s 2.26E-05 20 11.280; Surfill_w 4.51E-04; Surfill_z 4.51E-04; Surfsm_s 1.34E-05 39 6.720; Surfsm_w 2.69E-04; Surfsm_z 2.69E-04

	EXCHANGE 1
	-equilibrate 1
	X 1.80E-03 #illite
	Z 5.85E-03 #smectite
	-equilibrate 1
	SAVE SOLUTION 1
	SAVE SURFACE 1
	SAVE EXCHANGE 1
	END
	USE SOLUTION 1
	USE SURFACE 1
	USE EXCHANGE 1
	EQUILIBRIUM_PHASES 1
	Fix_pH -7.01 NaOH 1
	REACTION 1
	UO2Cl2
		U")
	100 SAVE fraction_U_sorbed
	-end
	SOLUTION 1 COX porewater 473-475m
	temp	80
	pH	7.01
	pe	3
	redox	pe
	units	mmol/l
	density	1
	Alkalinity 1.4 meq/l #2.2 meq/L #
	Ca	7.9
	Cl	40.4
	K	1.1
	Mg	5.3
	Na	42.7 charge
	S(6)	13.8
	Sr	0.2
	-water	1 # kg
	SAVE SOLUTION 1
	END
	USE SOLUTION 1
	SURFACE 1
	#name; -no_edl
	-equilibrate 1

i Concentration set up to 0.087 due to an experimental error. This has been taken into account in the modelling and calculations.

i Calculated considering the uncertainty of the alkalinity measurements. Maximum and minimum values of pCO2 calculated.

i Calculations in Phreeqc from the equilibrium between pH and alkalinity measured after the equilibrium with illite; maximum and minimum pCO2 considering the uncertainty of alkalinity measurements

Acknowledgements

I am not a person of many writing words, I spent most of my time talking and all my thanks were given to each one of you in particular and by person. Although, I want to leave it registered here. Most probably I will forget someone, but no hard feelings on me please! It's already quite difficult to find the right words to describe what each one of you have contributed for this period of my life.

Conclusions

Internal standard solution:

The internal standard solution is normally prepared to verify the signal of the measurements and to correct the drift of the analysis. Here, 205 Tl was used as internal standard. A solution of 2% ultrapure HNO3 and 2.5 ppb of 205 Tl was prior prepared for each dilution, such for the calibration curve.

The solution was firstly prepared by dilute 62.4 µL of 1002 µg/mL ICP-MS standard solution in 25 mL of 2% ultrapure HNO3 to obtain a solution of 25 ppm of 205 Tl. A second dilution was prepared by diluting 2 mL of 25 ppm 205 Tl solution in 2000 mL of 2% ultrapure HNO3.

Internal standard correction:

The internal standard signal was used to correct of drift from the matrix effect. The calculations were as follow:

Calibration curve:

The calibration curve was determined previously to each ICP-MS analysis. The 238 U standard solution of 999 µg/mL was used to prepare the standard solutions. Firstly, a mother solution of 500 ppb was prepared by diluting 25 µL of 238 U standard solution in 50 mL of 2% HNO3 + 205 Tl solution. 6 standard solutions were prepared as followed: The amount of exchange sites of the resin Dowex 1x2 can be calculated based on:

-Assuming 70 %wt is water retention capacity of the resin, 30 % is the dry fraction of the resin;

-The density of the resin (ρresin) is 0.705 g/cm3;

-The capacity of exchanging (CE) of 0.7 meq/mL by wetted bed volume;

B.2. Calculation of the sorption parameters of clay minerals

The main sorption parameters of the clay minerals taken into account in this study are: the grams of mass of each clay mineral, the respective number of surface and exchange sites and their specific surface. The specific surface was provided from the literature. The mass of clay mineral and number of sites were calculated as follow: The results are detailed in Table C. 1 and Table C. 2 and show that the most suitable procedure for transfer the samples from the incubator into the glove-box is using a pre-heated isolation box (internal and external layers) at 90 ºC using a faster closing lid system (Figure C. 2 and Figure C. 3). The system was validated in real conditions, where the increase of + 4ºC in the begging of the transfer and loss 1 ºC after 20 minutes. The temperature is maintained ± 5 °C for more than 20 minutes outside of the incubator. This time is enough to transport the samples and for the sampling. The

C.1.3. Tightness control test:

The possible leak of the tubes PPCO in temperature was tested by measure the volume of water before and after the agitation of the tubes along the time. Different sealing tapes were used: a) tubes closed with no sealing tape; b) tubes closed and sealed with parafilm; and c) tubes closed and sealed with Teflon tape. All experimental tests were performed at 80ºC.

The results detailed in Table C. 4 show that the absence of any sealing tape, after 6 days of the agitation, there is a loss of approximately 1% of the volume of the sample (Figure C. 6). For the test using the parafilm, the results show that the parafilm is melted, although using Teflon tape no significant losses of the solution were observed.

ternary species

The Ca-U(VI)-CO3 complexes have being studied in different experimental conditions (Bernhard et al., 2001;Dong and Brooks, 2006;Endrizzi and Rao, 2014;Kalmykov and Choppin, 2000;Lee and Yun, 2013). In Table D. 1, a summary of the experimental conditions considered for each work and the thermodynamic data for the Reaction 1 and Reaction 2 (see Chapter 2) are presented.

Calculations of the experimental conditions from the literature data were performed using Phreeqc with ThermoChimie database taking in account the ionic strength corrections (Davies or SIT conventions) (Giffaut et al., 2014). The results of the speciation of U(VI) as a function of Ca in solution and the saturation indexes of the most relevant minerals for each work are indicated in the following Figures (D. It should be noted Dong and Brooks (2006) data was already shown and discussed in Chapter 2, therefore we will not be shown here.

Appendix E: Thermodynamic data

In this appendix the thermodynamic data included in ThermoChimie database (Giffaut et al., 2014) for Ca-UO2 2+ -CO3 2-Na + -NO3 --H2O systems are shown in Table E. 10 and E. 11. This thermodynamic data is useful for the experiments of the ionic exchange and sorption of U(VI) on Illite, Callovo-Oxfordian clay fraction and claystone samples. All the geochemical model calculations were performed taking in account these thermodynamic data.