actinide sorption by cementitious materials · actinide/lanthanide sorption by cementitious...

PSI Bericht Nr. 18-02 August 2018

ISSN 1019-0643

Jan Tits and Erich Wieland

Paul Scherrer Institut (PSI) Nuclear Energy and Safety Research Department (NES)

Laboratory for Waste Management (LES) Forschungsstrasse 111

5232 Villigen PSI, SwitzerlandTel. +41 56 310 21 11Fax +41 56 310 21 99

www.psi.ch

Actinide Sorption by Cementitious Materials

Nuclear Energy and Safety Research Department (NES)Laboratory for Waste Management (LES)

PSI Bericht 18-02 2

Preface The Laboratory for Waste Management of the Nuclear Energy and Safety Research Department of the Paul Scherrer Institut is performing work to develop and test models as well as to acquire specific data relevant to performance assessments of planned Swiss nuclear waste repositories. These investigations are undertaken in close co-operation with, and with the partial financial support of, the National Cooperative for the Disposal of Radioactive Waste (Nagra).

I PSI Bericht 18-02

Abstract

The selection of well-supported sorption values for cement sorption databases for the assessment of the safety of a deep geological repository for low and intermediate level radioactive waste requires a detailed understanding of the radionuclide – cement interactions, the sorption mechanisms, and how radionuclide retention is affected by chemical conditions prevailing in the repository, such as pH and pore water composition. In the past years, a comprehensive research programme has been carried out by the Laboratory for Waste Management at the Paul Scherrer Institute with the aim to understand the interaction processes of actinides and lanthanides with cementitious materials both on a microscopic and on a macroscopic scale. Several spectroscopic studies that had been carried out at the Laboratory for Waste Management and that have been published earlier in the open literature reveal that incorporation into cement minerals is an important retention mechanism for actinides along with adsorption onto cement mineral surfaces, and further that calcium silicate hydrate (C–S–H) phases are the main uptake-controlling phase for actinides and lanthanides in cementitious materials. The present report contains an overview of the batch sorption studies with actinides and lanthanides in different redox states and on various cementitious materials carried out at the Laboratory for Waste Management. The actinides and lanthanides investigated include Eu(III) and Am(III), Th(IV), Np(IV), Np(V) and Np(VI) and U(VI). The experimental results and interpretations provided in this report aim to 1) quantify the actinide/lanthanide uptake by hardened cement paste (HCP) and C–S–H phases, a major component of HCP, 2) to describe the effects of cement paste / C–S–H composition, cement pore water composition and pH on actinide/lanthanide sorption by cementitious materials, and 3) to discuss the mechanisms driving the uptake of actinides by cement paste and CSH phases.

PSI Bericht 18-02 II

Zusammenfassung

Die Auswahl gut abgestützter Sorptionswerte für Zementsorptionsdatenbanken zur Beurteilung der Sicherheit eines geologischen Tiefenlagers für schwach- und mittelradioaktiver Abfälle erfordert ein detailliertes Verständnis der Radionuklid-Zement-Wechselwirkungen, der Sorptionsmechanismen und wie die Radionuklidretention durch die im Endlager vorherrschenden chemischen Bedingungen wie pH Wert und Porenwasserzusammensetzung beeinflusst werden. In den vergangenen Jahren wurde am Labor für Endlagersicherheit (LES) des Paul Scherrer Institut (PSI) ein umfangreiches Forschungsprogramm durchgeführt, mit dem Ziel die Prozesse der Wechselwirkung von Aktiniden und Lanthaniden mit zementartigen Materialien sowohl im mikroskopischen als auch im makroskopischen Massstab zu verstehen. Mehrere spektroskopische Studien, die am LES durchgeführt worden waren und veröffentlicht sind, zeigen, dass der Einbau in Zementmineralien, nebst der Adsorption an Zementmineraloberflächen, ein wichtiger Retentionsmechanismus für die Aktiniden ist und dass Calciumsilikathydrat (C-S-H)-Phasen die wichtigsten Phasen bei der Kontrolle der Retention von Aktiniden und Lanthaniden durch Zementstein sind. Der vorliegende Bericht gibt einen Überblick über die am LES durchgeführten Batch-Sorptionsstudien mit Aktiniden und Lanthaniden in unterschiedlichen Redoxzuständen und auf verschiedene zementartige Materialien. Die untersuchten Aktiniden und Lanthaniden umfassen Eu(III) und Am(III), Th(IV), Np(IV), Np(V) und Np(VI) und U(VI). Die experimentellen Ergebnisse und deren Interpretationen in diesem Bericht zielen darauf ab 1) die Aktiniden/Lanthaniden-Sorption durch Zementstein und C-S-H Phasen, die wichtigste Komponente von Zementstein, zu quantifizieren, 2) die Auswirkungen der Zement/C-S-H Zusammensetzung, der Zementporenwasserzusammensetzung und des pH Wertes auf die Aktiniden/Lanthaniden-Sorption durch zementartige Materialien zu beschreiben, und 3) die Mechanismen zu diskutieren, welche die Aufnahme von Aktiniden durch Zementstein und C-S-H Phasen bewirken.

III PSI Bericht 18-02

Résumé

La sélection de valeurs de sorption fiables pour les bases de données sur la sorption du ciment pour l'évaluation de la sûreté des sites de stockage de déchets faiblement et moyennement radioactifs en couche géologique profonde exige une compréhension détaillée des interactions radionucléides-ciment, des mécanismes de sorption et de la façon dont la rétention des radionucléides est affectée par les conditions physico-chimiques régnants dans l’ouvrage de stockage, comme le pH et la composition de l'eau interstitielle. Les dernières années, des efforts importants ont été déployés au Laboratoire Sûreté des dépôts de déchets radioactifs de l'Institut Paul Scherrer pour élucider les processus contrôlant l'adsorption d'actinides et de lanthanides par les matériaux cimentaires à la fois à l'échelle microscopique et à l'échelle macroscopique. Plusieurs études spectroscopiques qui ont été réalisées au Laboratoire Sûreté des dépôts de déchets radioactifs et qui ont été publiées dans le passé dans la littérature ouverte, révèlent que l'incorporation dans les minéraux du ciment ainsi que l'adsorption sur les surfaces minérales du ciment jouent un rôle important dans la rétention des actinides, et que les phases de silicate de calcium hydraté (phases C-S-H) sont les principales phases contrôlant l'absorption des actinides et des lanthanides dans les matériaux cimentaires. Le présent rapport donne un aperçu des études de sorption de batch avec des actinides et des lanthanides dans différents états redox et des matériaux cimentaires réalisées au Laboratoire Sûreté des dépôts de déchets radioactifs. Parmi les actinides et lanthanides étudiés, comptent Eu(III) et Am(III), Th(IV), Np(IV), Np(IV), Np(V) et Np(VI) et U(VI). Les résultats expérimentaux fournis dans ce rapport visent à 1) quantifier la rétention des actinides/lanthanides par la pâte de ciment durcie (HCP) et les phases C-S-H, un composant majeur de HCP, 2) décrire les effets de la composition de la pâte de ciment / C-S-H, ainsi que de la composition et du pH de la solution interstitielle sur la sorption des actinides/lanthanides par les matériaux cimentaires, et 3) discuter les mécanismes d'adsorption d'actinides par le la pâte de ciment et les phases C-S-H.

PSI Bericht 18-02 IV

Table of contents

Abstract ................................................................................................................................... I

Zusammenfassung ...................................................................................................................... II

Résumé ................................................................................................................................ III

Table of contents ....................................................................................................................... IV

Tables .............................................................................................................................. VII

Figures ............................................................................................................................. VIII

1 Introduction ............................................................................................................ 1 1.1 Cementitious materials as barrier for actinides in low and intermediate level

radioactive waste repositories ................................................................................... 1 1.2 Aim and structure of this report ................................................................................ 1

2 Chemistry of cementitious materials and actinides/lanthanides ........................ 3 2.1 Cementitious materials ............................................................................................. 3 2.1.1 Cement paste ............................................................................................................. 3 2.1.2 C–S–H phases .......................................................................................................... 4 2.2 Actinide chemistry .................................................................................................... 7 2.3 State of the art on actinide sorption on cement minerals .......................................... 8 2.4 Description of the experimental studies.................................................................... 9

3 Materials and methods ......................................................................................... 13 3.1 Materials ................................................................................................................. 13 3.1.1 Experimental cement pore waters ........................................................................... 13 3.1.2 Titanium dioxide ..................................................................................................... 13 3.1.3 C–S–H phases ......................................................................................................... 14 3.1.4 Hardened cement paste ........................................................................................... 14 3.2 Analytical methods ................................................................................................. 15 3.2.1 Radiotracers and radio-assay .................................................................................. 15 3.2.2 Chemical analyses .................................................................................................. 15 3.2.3 Experimental uncertainties and reproducibility ...................................................... 16 3.3 Experiments ............................................................................................................ 16 3.3.1 Characterization of the C–S–H phases and HCP material ...................................... 16 3.3.2 Recrystallization experiments with C–S–H phases ............................................... 17 3.3.3 Preliminary solubility tests of actinide/lanthanide solutions .................................. 17 3.3.4 Sorption experiments .............................................................................................. 19 3.3.5 Co-precipitation experiments with C–S–H phases ................................................ 19 3.3.6 Desorption experiments on C–S–H phases ............................................................ 20 3.4 Interpretation of the sorption data ........................................................................... 20

V PSI Bericht 18-02

3.4.1 Distribution ratios and concentration calculations .................................................. 20 3.4.2 Modelling the effect of hydrolysis on sorption ....................................................... 21

4 Characterization of the C–S–H phases .............................................................. 23 4.1 Pre-equilibration of solid phases ............................................................................ 23 4.2 Solid phase and aqueous phase composition .......................................................... 24 4.3 45Ca uptake by C–S–H phases ................................................................................ 25

5 Sorption of trivalent actinides and lanthanides ................................................. 28 5.1 Speciation ............................................................................................................... 28 5.2 Preliminary solubility experiments ......................................................................... 30 5.3 Eu(III) sorption and co-precipitation kinetics ......................................................... 31 5.3.1 C–S–H phases ........................................................................................................ 31 5.3.2 Hardened cement paste ........................................................................................... 33 5.4 Desorption tests ...................................................................................................... 33 5.5 Effect of the S:L ratio ............................................................................................. 34 5.6 Sorption isotherms .................................................................................................. 35 5.7 Discussion ............................................................................................................... 37 5.7.1 Sorption data ........................................................................................................... 37 5.7.2 Uptake mechanisms ................................................................................................ 38

6 Sorption of tetravalent actinides ......................................................................... 40 6.1 Speciation ............................................................................................................... 40 6.2 Preliminary solubility tests ..................................................................................... 42 6.3 Sorption and co-precipitation kinetics .................................................................... 43 6.3.1 Sorption kinetics ..................................................................................................... 43 6.3.2 Th(IV) co-precipitation kinetics ............................................................................. 45 6.4 Desorption experiments .......................................................................................... 46 6.5 Effect of S:L ratio ................................................................................................... 47 6.6 Sorption isotherms .................................................................................................. 47 6.7 Effect of the C–S–H composition (C:S ratio) ......................................................... 49 6.8 Discussion ............................................................................................................... 49 6.8.1 Sorption data ........................................................................................................... 49 6.8.2 Uptake mechanisms ................................................................................................ 50

7 Sorption of pentavalent actinides ........................................................................ 51 7.1 Speciation ............................................................................................................... 51 7.2 Preliminary solubility tests ..................................................................................... 53 7.3 Sorption kinetics ..................................................................................................... 55 7.4 Desorption tests ...................................................................................................... 56 7.5 Sorption isotherms .................................................................................................. 57 7.6 Effect of pH and aqueous Ca concentration on the Np(V) sorption onto TiO2 ...... 58 7.7 Effect of the C:S ratio of C–S–H phases ............................................................... 59 7.8 Discussion ............................................................................................................... 60

PSI Bericht 18-02 VI

7.8.1 Sorption data ........................................................................................................... 60 7.8.2 Uptake mechanisms ................................................................................................ 63

8 Sorption of hexavalent actinides .......................................................................... 64 8.1 Speciation ............................................................................................................... 64 8.2 Preliminary solubility tests ..................................................................................... 66 8.3 Sorption kinetics ..................................................................................................... 68 8.4 Effect of the S:L ratio ............................................................................................. 70 8.5 Sorption isotherms .................................................................................................. 71 8.6 Effect of pH and aqueous Ca concentration on the U(VI) and Np(VI)

sorption onto TiO2 .................................................................................................. 72 8.7 Effect of the C:S ratio ............................................................................................. 74 8.8 Discussion ............................................................................................................... 76 8.8.1 Sorption data ........................................................................................................... 76 8.8.2 Uptake mechanisms ................................................................................................ 77

9 Summary and conclusions ................................................................................... 78

10 Acknowledgements ............................................................................................... 79

11 References .............................................................................................................. 79

VII PSI Bericht 18-02

Tables

Table 1: Bulk composition of selected C–S–H phases with different C:S ratios synthesized in ACW at pH = 13.3 and in MilliQ water .......................................... 24

Table 2: Relevant thermodynamic complexation constants for trivalent actinides and lanthanides used in the speciation calculations presented in Fig. 6 (Hummel et al., 2002; Guillaumont et al., 2003; Neck et al., 2009; Thoenen et al., 2014) ....................................................................................................................... 29

Table 3 Equilibrium constants for An(IV) aqueous hydroxide complexes, ternary Ca-An(IV)-OH complexes and solid An(IV) hydroxides (Guillaumont et al., 2003; Altmaier et al., 2008; Rai et al., 2008; Rand et al., 2008; Fellhauer et al., 2010; Thoenen et al., 2014). ............................................................................. 41

Table 4: Summary of log10*β° and log10*K°s,0 used in the Np(V) speciation

calculations. Only species formed between 10 < pH < 14 are considered. (Guillaumont et al., 2003; Altmaier et al., 2013; Thoenen et al., 2014; Fellhauer et al., 2016a) ........................................................................................... 52

Table 5: Overall and stepwise formation constants for relevant Np(V) hydrolysis complexes used in the model calculations (Log*Kn

0 = Log*βn0 - Log*βn-1). ......... 59

Table 6: Stepwise hydrolysis constants for different actinides and lanthanides at pH = 10. Comparison with the log Rd values of the respective actinides and lanthanides. ............................................................................................................. 63

Table 7: Summary of log10*β° and log10*K°s,0 used in the An(VI) speciation calculations (An(VI) = U(VI) or Np(VI)). Only species formed between 10 < pH < 14 are considered (Gaona et al. (2013a), Thoenen et al. (2014), Altmaier et al. (2017)). ........................................................................................... 64

Table 8: Overall and stepwise formation constants for relevant Np(VI) hydrolysis complexes used in the model calculations. (log*Kn

0 = log*βn0 - log*βn-1). ............. 73

PSI Bericht 18-02 VIII

Figures

Fig. 1: Degradation of hydrated cement by leaching with ground water. a) Evolution of the hydrated cement composition with progressing degradation. b) Evolution of the cement pore water as function of the volume of leach water (modified from Jacques et al., 2010) ........................................................................ 3

Fig. 2a: Schematic of the tobermorite-based structure of a C–S–H phase with a C:S ratio of 0.66 and a maximum degree of protonation of the silica chains (adapted from Richardson, 2008). ............................................................................ 5

Fig. 2b: Schematic of the tobermorite-based structure of a C–S–H phase with a C:S ratio of 1.5 obtained by removal of the bridging tetrahedra and neutralising the negative charges of the silanol groups by Ca2+ ions (adapted from Richardson, 2008). .................................................................................................... 5

Fig. 3: Evolution of the aqueous phase composition during the synthesis of C–S–H phases with three different C:S ratios in ACW at pH = 13.3. S:L = 0.02 kg L-

1. a) Ca concentration. b) Si concentration ............................................................. 23

Fig. 4: Ca (a) and Si(b) solubility of C–S–H phases as a function of the actual C:S ratio in alkali-free solutions and in ACW at a constant pH of 13.3. Each datapoint represents an individual C–S–H phase. C:S ratios were calculated from mass balance calculations. The effective C:S ratio of the C–S–H phases with target C:S ratios ≥ 1.5 (in alkali-free solution) and ≥ 1.2 (in ACW) was corrected for the portlandite content based on the data in Table 1. ........................ 25

Fig. 5: 45Ca activity ratios for two C–S–H recrystallization experiments (S:L ratio = 2·10-3 kg L-1) in a) alkali-free solution and b) in ACW. Red lines in (a) and (b) are fits with a model combining homogeneous incorporation with surface sorption. c) fraction of C–S–H recrystallized ......................................................... 27

Fig. 6: Trivalent actinide and lanthanide speciation as a function of pH in the absence of carbonates. The total Am(III) and Eu(III) concentrations were fixed at 10-5 M. a and b) Am(III) speciation (and Nd(III) speciation) in absence of Ca and in the presence of 2·10-2 M Ca, c) Eu(III) speciation. Thermodynamic equilibrium constants taken from Table 2. .................................. 30

Fig. 7: Eu(III) concentration measured in solution before and after centrifugation as a function of the input Eu(III) concentration. Measurements were conducted after 1 day (a, b) and 30 days equilibration (c, d). The solid lines in b and d represent 100% Eu(III) recovery in solution. The dashed horizontal lines represent the calculated Eu(OH)3(am) solubility limit in ACW based on the thermodynamic data for Nd(III). Dotted lines are added to guide the eye. ............ 31

Fig. 8: a) Eu(III) sorption kinetics onto C–S–H phases with different C:S ratios in ACW at pH = 13.3. Experimental conditions: S:L = 5·10-4 kg L-1, 1.98·10-9 M < [Eu]tot < 9.9·10-9 M. Red symbols: experiments carried out with dried C–S–H phases. Black symbols: experiments carried out with fresh C–S–H suspensions. b) Eu(III) – C–S–H co-precipitation kinetics in ACW at pH = 13.3. Comparison with the data from the sorption kinetic experiments (datapoints in grey). The dashed lines represent the Rd,max value. ......................... 32

Fig. 9: Eu(III) sorption kinetics (a) and Am(III) sorption kinetics (b) onto HCP in ACWHCP-I (pH = 13.3) and ACWHCP-II (pH = 12.5). Experimental conditions: S:L = 10-4 kg L-1, [Eu]tot = 2·10-9 M, [Am]tot = 3·10-10 M. ...................................... 33

IX PSI Bericht 18-02

Fig. 10: Eu(III) desorption tests onto C–S–H phases in ACW at pH = 13.3 after one day and 60 days sorption. a) C:S = 0.75; b) C:S = 1.25. Experimental conditions: S:L = 2.5·10-3 kg L-1, [Eu]tot = 3.0·10-8 M. The dashed lines represent the Rd,max value. ....................................................................................... 34

Fig. 11: Eu(III) sorption onto C-S-H phases and HCP as function of the S:L ratio. a) C–S–H phase (C:S = 1.07) in ACW at pH = 13.3. Experimental conditions: equilibration time = 2 weeks, [Eu]tot = 8.5·10-10 M. The dashed line corresponds to the Rd,max values as function of the S:L ratio. b) Fresh HCP in ACW-I at pH 13.3: equilibration time = 2 weeks, [Eu]tot = 7.2·10-10 M. ............................................................................................................................ 35

Fig. 12 a,b) Eu(III) sorption isotherm in ACW onto a C–S–H phase with C:S ratio = 1.0. Experimental conditions: S:L = 5·10-4 kg L-1, equilibration time = 2 weeks. Cm(III) sorption experiments: S:L = 10-3 kg L-1. c,d) Eu(III) and Am(III) sorption isotherms onto fresh HCP. Experimental conditions: S:L = 10-4 kg L-1; equilibration time = 2 weeks. The dashed line in a), represents Rd,max values for Eu(III). ......................................................................................... 36

Fig. 13: Fluorescence emission spectra of 10-7 M Cm(III) sorbed onto a C–S–H phase (C:S = 1.07) in ACW at pH 13.3. a) Fluorescence emission as function of the reaction time. b) Fluorescence emission spectra after 58 days contact time recorded at increasing delay times. The time dependence of the emission decay shows a bi-exponential decay behaviour. A detailed discussion of the Cm(III) TRLFS experiments is given in the paper of Tits et al. (2003). Reprinted with permission from Tits et al. (2003). Copyright 2005 by Walter de Gruyter GmbH. ................................................................................. 39

Fig. 14: Tetravalent actinide speciation as a function of pH in the absence of carbonates. [Th(IV)]tot = 10-5 M. I = 0.3 M. a) [Ca] = 5·10-5 M and [Si(OH)4] = 8·10-3 M, b) [Ca] = 10-3 M and [Si(OH)4] = 10-3 M, c) [Ca] = 2·10-2 M and [Si(OH)4] = 10-4 M. Thermodynamic equilibrium constants for solid hydroxides and aqueous hydroxyl complexes are taken from the NEA thermodynamic databases (Guillaumont et al., 2003; Rand et al., 2008) and from Fellhauer et al., (2010). .................................................................................. 42

Fig. 15: Solubility tests with Th(IV) in ACW at pH = 13.3. a) Percentage of the Th(IV) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Th(IV) inventory. b) Th(IV) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Th(IV) inventory. Equilibration times were 1 day and 21 days. The curves in a) only serve to guide the eye. The shaded area in b) represents the calculated Th(IV) solubility limit with its uncertainty in ACW with respect to ThO2(am, fresh, hyd). The line with slope +1 represents 100% recovery in solution. ...................... 43

Fig. 16: Np(IV) and Th(IV) sorption kinetics on C–S–H phases having various compositions in alkali-free conditions and in ACW at pH = 13.3. a) Np(IV) sorption in alkali-free conditions (pH 10.1), b) Np(IV) sorption in ACW (pH 13.3), c) Th(IV) sorption in alkali-free conditions (pH 10.1), d) Th(IV) sorption in ACW (pH 13.3). Experimental conditions: Np(IV) sorption kinetics: S:L = 2·10-4 kg L-1, [Np(IV)]tot = 1.7·10-7 M (237Np + 239Np) and 1.3·10-10 M (239Np). Th(IV) sorption kinetics: S:L = 2·10-3 kg L-1, [Th(IV)]tot = 2·10-11 M. ............................................................................................................. 44

PSI Bericht 18-02 X

Fig. 17: Th(IV) sorption kinetics onto fresh HCP in ACWHCP-I (pH = 13.3) and onto aged HCP (ACWHCP-II, pH = 12.5). Experimental conditions: S:L = 1.0·10-4 kg L-1, [Th(IV)]tot = 2·10-10 M. ................................................................................ 45

Fig. 18: Th(IV) co-precipitation kinetics onto C–S–H phases with various C:S ratios in a) alkali-free systems and b) ACW (pH 13.3). Experimental conditions: S:L = 2·10-3 kg L-1, [Th]tot = 2·10-11 M. Comparison with the data from the sorption kinetic experiments (datapoints in grey) ................................................... 46

Fig. 19: Th(IV) sorption (a) and co-precipitation (b) experiments on a C–S–H phase with a C:S ratio of 1.07 in ACW at pH 13.3. Desorption tests conducted after 1 day, 20 day and 120 days of sorption/co-precipitation. Experimental conditions: S:L = 2·10-3 kg L-1, [Th]tot = 2·10-12 M, desorption time = 3 days. ...... 46

Fig. 20: Th(IV) sorption as a function of the S:L ratio in ACW at pH = 13.3 on C–S–H phases with different C:S ratios (a) and on fresh HCP (b). Experimental conditions for the C–S–H systems: [Th]tot = 5·10-10 M, equilibration time = 2 weeks. Experimental conditions for fresh HCP in ACWHCP-I: [Th]tot = 7·10-10 M, equilibration time = 2 weeks. ............................................................................ 47

Fig. 21: a, b) Th(IV) sorption isotherm onto a C–S–H phase with C:S = 1.07. Experimental conditions: S:L = 5·10-4 kg L-1, equilibration time = 2 weeks. c, d) Th(IV) sorption isotherm onto fresh HCP in ACWHCP-1 (pH = 13.3) and in aged HCP in ACWHCP–II (portlandite saturated solution, pH = 12.5). Experimental conditions: S:L = 10-4 kg L-1, equilibration time = 2 weeks. Black vertical lines: Experimental Th(IV) stability limit. Red vertical lines: Thermodynamic Th(IV) solubility limit. ................................................................ 48

Fig. 22: An(IV) sorption onto C–S–H phases as function of the C:S ratio in alkali-free solutions and in ACW (pH 13.3). Experimental conditions for Np(IV) experiments: [Np]tot = 2·10-10 M, S:L = 5·10-3 kg L-1, equilibration time = 3 days. Experimental conditions for the Th(IV) experiments: [Th]tot = 8.5·10-9 M, S:L = 5·10-3 kg L-1, equilibration time = 3 days. ............................................... 49

Fig. 23: Np(V) speciation calculations in the pH range 10 < pH < 14 in the absence of Ca (a) and in the presence of low (b) and high (c) Ca concentrations representing the typical Ca concentrations in C–S–H solutions in the absence of alkalis. ................................................................................................................ 52

Fig. 24: Solubility tests of Np(V) in a) 10-4 M Ca(OH)2 at pH = 10.3, b) 10-2 M Ca(OH)2 at pH = 12.3, and c) 0.3 M NaOH and ACW. Np(V) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Np(V) inventory. Equilibration time was 30 days. The line with slope +1 represents 100% recovery in solution. ......... 54

Fig. 25: Np(V) sorption kinetics on C–S–H phases with three different C:S ratios in alkali-free conditions and on fresh HCP at pH 13.3. a) Aqueous Np(V) equilibrium concentrations versus time. b) Rd values versus time. Experimental conditions for the C–S–H experiments: S:L = 2·10-4 kg L-1, [Np]tot = 3.7·10-6 M, 2·10-8 M < [Np]eq < 10-7 M. Experimental conditions for the fresh HCP experiments in ACWHCP-I: S:L = 10-4 kg L-1, [Np]tot = 5·10-6 M, 2·10-7 M < [Np]eq < 5·10-7 M. ........................................................................... 55

Fig. 26: Np(V) desorption kinetics after 1 day sorption and 30 days or 60 days sorption. a) C–S–H phase with C:S ratio = 0.65 in alkali-free solution. b) fresh HCP in ACWHCP-I at pH = 13.3. Experimental conditions for the C–S–H system: S:L = 1.25·10-4 kg L-1, [Np]tot = 3.6·10-5 M, 10-7 M < [Np]eq <

XI PSI Bericht 18-02

4·10-7 M. Rd,max = 7·108 L kg-1. Experimental conditions for the HCP system: S:L = 1.25·10-4 kg L-1, [Np]tot = 3.6·10-5 M, 7·10-7 M < [Np]eq < 10-6 M. .............. 56

Fig. 27: Np(V) sorption isotherms onto C–S–H phases with varying C:S ratios in alkali-free solution and in ACW at pH = 13.3. a, c) The amount of Np(V) sorbed versus the Np(V) aqueous equilibrium concentration. b, d) Rd(Np(V) versus the Np(V) aqueous equilibrium concentration. Experimental conditions: S:L = 5·10-3 kg L-1 (exp. 1) and 2·10-4 kg L-1 (exp. 2), equilibration time = 2 weeks. The red and black dashed vertical lines represent the experimental solubility limits determined previously (Fig. 24). The Rd,max values are shown as black dotted lines (Figs. b and d). ......................... 57

Fig. 28: Np(V) sorption onto TiO2. Effect of pHc (a) and of the aqueous Ca2+ concentration in ACW (b). The solid line in (a) represents the model calculation using the equation Rd = Rd

0/Fred. The solid line in (b) is added to guide the eye. .......................................................................................................... 59

Fig. 29: Sorption of Np(V) on C–S–H phases in alkali-free conditions (10 < pH < 12.5) and in ACW at a constant pH of 13.3. a) Effect of the C:S ratio in alkali-free conditions. b) Rd values from a) plotted versus pHc. c) Rd values from a) measured in ACW at fixed pH = 13.3 plotted versus the free Ca2+ concentration .......................................................................................................... 61

Fig. 30: U(VI) speciation calculations in the pH range 10 < pH < 14 in the absence of Ca (a) and in the presence of high Ca concentrations (b) representing the maximal Ca concentration in C–S–H solutions in the absence of alkalis. .............. 65

Fig. 31: Solubility tests with U(VI) in alkali-free solutions. Figs. a, c, e: Percentage of the U(VI) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Figs. b, d, f: U(VI) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Equilibration time was 1 day and 7 days in alkali-free systems. The curves in Figs. a, c, e only serve to guide the eye. The lines with slope +1 in Figs. b, d, f represent 100% recovery in solution. ....................... 67

Fig. 32: Solubility tests with U(VI) in ACW. a) Percentage of the U(VI) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. b) U(VI) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Equilibration time was 1 day and 32 days. The curves in Fig. a only serve to guide the eye. The lines with slope +1 in Fig. b represent 100% recovery in solution. .................................................................................................................. 68

Fig. 33: Sorption kinetics (a) and co-precipitation kinetics (b) of U(VI) by C–S–H phases with different C:S ratios in alkali-free solutions (pH between 10.0 and 12.5) and in ACW (pH 13.3). Experimental conditions: S:L ratio = 5·10-

3 kg L-1, [U]tot = 5.4·10-7 M, Rd,max = 1.5·106 L kg-1. ............................................. 69

Fig. 34: Sorption kinetics of U(VI) onto fresh HCP in ACWHCP-I (pH 13.3). Experimental conditions: S:L = 10-3 kg L-1, [U]tot = 5.3·10-7 M. ............................ 70

Fig. 35: U(VI) sorption as a function of the S:L ratio in ACWHCP-I at pH = 13.3 on fresh HCP. Experimental conditions: [U(VI)]tot = 5.4·10-7 M; equilibration time = 30 days. ........................................................................................................ 70

PSI Bericht 18-02 XII

Fig. 36: U(VI) sorption isotherms on C–S–H phases with different Ca:Si ratios in alkali-free conditions and in ACW at pH = 13.3. a) The amount of U(VI) sorbed versus the U(VI) aqueous equilibrium concentration. b) Rd versus the U(VI) aqueous equilibrium concentration. Vertical grey bars represent the solubility of Na–uranate in the ACW. Dotted lines represent the Rd,max values. Dashed and solid lines are added to guide the eye. .................................... 71

Fig. 37: U(VI) sorption isotherms on fresh HCP in ACWHCP-I (pH 13.3). a) The amount of U(VI) sorbed versus the aqueous equilibrium concentration. b) Rd versus the aqueous equilibrium concentration. Thermodynamic U(VI) solubility taken from Fig. 30b. Experimental stability limit taken from Fig. 32b. ......................................................................................................................... 72

Fig. 38: Sorption of Np(VI) on TiO2. a) Effect of pH in absence of Ca. b) Effect of the Ca2+ concentration in 0.3 M NaOH. Experimental data (symbols) and modeling (line in Fig. a) (see text). Rd, max = 4.3·107 L kg-1. The line in Fig. 38b is added to guide the eye. ................................................................................. 74

Fig. 39: Uptake of U(VI) and Np(VI) by C–S–H phases under alkali-free conditions (10 < pH < 12.5) and in ACW at a fixed pH of 13.3. a) Effect of the C:S ratio, b) Rd values from a) plotted versus pHc, c) Rd values determined in ACW at a constant pH = 13.3, and plotted versus the aqueous Ca2+ concentration. Rd, max = 4.3·107 L kg-1 (239Np experiment) and 1.5·107 L kg-1 (233U experiment). ................................................................................................... 75

1 PSI Bericht 18-02

1 Introduction

1.1 Cementitious materials as barrier for actinides in low and intermediate

level radioactive waste repositories

Permanent storage in stable deep geological formations is accepted as the preferred solution for the disposal of radioactive waste in many countries. Such a disposal concept includes a combination of engineered and natural barriers aiming at immobilizing the radionuclides to prevent their release into the environment. In the case of low and intermediate level radioactive waste (L/ILW), cementitious materials are widely used for the construction of the engineered barrier system. In addition to their low cost and their simple implementation, these materials are attractive owing to their many favourable chemical properties w.r.t. the immobilization of radioactive contaminants: i) They are capable of maintaining a high pH (10 < pH < 13.5) for a long time, thus limiting the solubility of many radionuclides, and ii) they are known to possess a high sorption capacity for many ions, thus retarding the migration of radionuclides from the waste containers to the host rock surrounding the repository (e.g. Chen et al., 2009 and references therein). The high sorption capacity of hydrated cementitious materials originates primarily from the presence of large amounts of quasi amorphous calcium silicate hydrate (C–S–H ) phases having very high reactive surface areas thus providing large numbers of surface sites accessible to metal cations present in radioactive waste for surface complexation. In addition, radionuclides may become incorporated upon recrystallization of the C–S–H phases. Incorporation may also take place when pore waters from the rock surrounding the repository, having a near-neutral pH, migrate into the cementitious near-field thus triggering dissolution re-precipitation processes of the C–S–H phases eventually leading to incorporation of radionuclides.

Actinides arising from resins, fuel cladding as well as from contaminated materials from reactor decommissioning, are an important class of safety-relevant radionuclides in L/ILW (Nagra, 2014). Lanthanides are often used as chemical analogues for the actinides with the same redox state in experimental studies because of their lower radiotoxicity.

1.2 Aim and structure of this report

During the last decade, major efforts have been made at the Paul Scherrer Institute to elucidate the processes controlling the uptake of actinides and lanthanides by cementitious materials on a microscopic scale. These studies revealed that incorporation into cement minerals is an important retention mechanism for actinides along with adsorption onto cement mineral surfaces, and that C–S–H phases are the main uptake-controlling phase for actinides and lanthanides in cementitious materials. Parallel to this research, a large effort was undertaken to evaluate the implications of these findings on the macroscopic sorption behaviour of actinides. A detailed understanding of the macroscopic actinide sorption behaviour is indispensable for the compilation of reliable sorption databases which are needed for the assessment of the safety of L/ILW repositories.

Macroscopic wet chemistry studies have been carried out onto hardened cement paste (HCP) as well as onto C–S–H phases, the cement component responsible for the favourable sorption properties of HCP. The studies included recrystallization studies of C–S–H phases as well as batch sorption experiments on C–S–H phases and HCP and in some cases, co-precipitation experiments, with C–S–H and trivalent actinides (Cm(III), Am(III)) and Eu(III) used as a

PSI Bericht 18-02 2

chemical analogue for trivalent actinides), tetravalent actinides (Th(IV), Np(IV)), pentavalent actinides (Np(V)) and hexavalent actinides (U(VI), Np(VI)). Sorption and desorption kinetics, sorption isotherms, the effect of S:L ratio and the influence of the C–S–H composition (C:S ratio) were investigated. In addition, for some actinides, reversibility of the sorption processes was explored. For pentavalent and hexavalent actinides, the sorption studies on cementitious materials were complemented with a sorption study on TiO2. Sorption studies on such very stable metal oxides with very low recrystallization rates give insight in the surface complexation behaviour of actinides under highly alkaline conditions. Differences between the sorption behaviour on TiO2 and on C–S–H phases were expected to allow incorporation into the C–S–H structure to be distinguished from surface complexation. Sorption experiments on TiO2 were carried out as a function of pH and aqueous Ca concentration, the two main variables in the cement pore water composition.

The present report provides a synthesis of the results from wet chemical sorption studies and summarizes hitherto unpublished wet chemistry and sorption data of actinides for use in safety assessments. The report is considered to be complementary to the spectroscopic studies previously published by the authors in the open literature

Chapter 2 provides some basics of the cement and actinide chemistry relevant for the understanding of actinide sorption by cementitious materials. Chapter 3 gives a detailed description of the cementitious materials used in the experiments, the experimental techniques, and the set-up of the different types of experiments. Chapter 4 describes the results of the solid and liquid phase characterization. The results of the actinide (An) and lanthanide (Ln) sorption experiments are described in four different chapters (chapters 5 to 8), one for each of the relevant redox states. Each of these chapters are subdivided in subsections dealing with the different parameters potentially influencing sorption as described above.

3 PSI Bericht 18-02

2 Chemistry of cementitious materials and

actinides/lanthanides

2.1 Cementitious materials

2.1.1 Cement paste Ordinary portland cement consists essentially of the clinker minerals: calcium silicates (Ca3SiO3 or C3S and Ca2SiO4 or C2S), tricalcium aluminate (Ca3Al2O6 or C3A) and tetracalcium alumino ferrite (Ca4(AlxFe(1-x)4O10) or C4AF)) together with minor amounts of other minerals such as gypsum, alkali sulphates, calcium carbonate, calcium oxide, magnesium oxide (Lothenbach and Wieland, 2006; Lothenbach and Winnefeld, 2006). During hydration, the cement components react with water and dissolve at different rates to form hardened cement paste (HCP) consisting mainly of four types of hydration products: calcium silicate hydrates (C–S–H phases), portlandite (calcium hydroxide), and to a minor extend AFm-phases (mainly calcium monocarboaluminate) and AFt phases (ettringite). C–S–H phases are the main reaction product resulting from the hydration process.

Infiltration of groundwater after closure of the repository induces the degradation of the cement paste mainly by sequential leaching of the cement phases. The degradation reaction includes dissolution of the original hydrated cement phases and precipitation of new phases (e.g. C–S–H phases with different composition, carbonates,…). It is well established that the cement degradation process comprises different stages (Jacques et al., 2010) (Fig. 1). The cement pore water in equilibrium with a fresh cement paste has a pH of ~13.3 and a composition dominated by NaOH and KOH, saturated with Ca(OH)2 (Fig. 1b; region I). With progressing degradation, the various cement components are sequentially leached. C–S–H phases of all cement phases are the most resistant to leaching.

a)

b)

Fig. 1: Degradation of hydrated cement by leaching with ground water. a) Evolution of the hydrated cement composition with progressing degradation. b) Evolution of the cement pore water as function of the volume of leach water (modified from Jacques et al., 2010)

PSI Bericht 18-02 4

The cement pore water evolves from a NaOH and KOH dominated solution to a Ca(OH)2 solution saturated w.r.t. portlandite and a pH of 12.5 (Fig. 1b; region II). After leaching the portlandite, the cement degradation reaches a third stage characterized by the incongruent dissolution of the C–S–H phases (Fig. 1b; region III). The C–S–H composition changes in this phase from a C:S ratio of ~1.7 to ~0.7 and the pH of the cement pore water decreases to a value of ~10.

In order to study the sorption behaviour of actinides on cementitious materials, C–S–H phases are selected as a representative cement mineral. Several macro- and micro spectroscopic studies revealed that the coordination environments of actinides sorbed on fresh and degraded HCP is very similar to that of actinides sorbed onto C–S–H phases (Stumpf et al., 2004; Mandaliev et al., 2009; Wieland et al., 2010; Macé et al., 2013). These observations provided ample evidence that C–S–H phases are the main sink for actinides in cementitious materials and that, in a cement environment, the same uptake processes take place as observed for synthetic C–S–H phases. These studies provided the justification for the choice of C–S–H phases as the main sorbent in the sorption studies with actinides.

2.1.2 C–S–H phases The structure of C–S–H phases is related to 14 Å tobermorite. They have a layered structure consisting of central Ca-O sheets with sevenfold coordinated calcium atoms, where all oxygen atoms are shared with silica atoms arranged in a “dreierketten” formation. The latter formation corresponds to silica dimers connected by bridging silica tetrahedra on both sides of the CaO2 sheets (e.g. Chen et al., 2004; Richardson, 2004; Bonaccorsi and Merlino, 2005; Garbev et al., 2008a; Garbev et al., 2008b; Richardson, 2008; Renaudin et al., 2009; Myers et al., 2013). The C:S ratio of the C–S–H phases can vary between 0.66 (Fig. 2a) and 1.5 (Fig. 2b). This is achieved basically by progressively replacing the interlayer protons neutralising the negative charges on the bridging tetrahedra by interlayer calcium atoms and by progressive removal of bridging tetrahedra from the silicate chains (e.g. Bonaccorsi et al., 2004; Nonat, 2004; Richardson, 2004; 2008; Lothenbach and Nonat, 2015). In fresh HCP, C–S–H phases with C:S ratios of 1.7 are commonly observed (Lothenbach and Nonat, 2015). Recent studies based on molecular modelling tend to support the idea that also at these high C:S ratios the C–S–H structure is tobermorite-like rather than a jennite-like structure as proposed in earlier studies (Nonat, 2004; Lothenbach and Nonat, 2015). The high C:S ratios are now believed to result from the accumulation of supplementary Ca in the C–S–H interlayer. Cement degradation caused by the intrusion of groundwater in the cementitious near-field causes incongruent dissolution of C–S–H phases, leading to lower C:S ratios.

The “dreierketten” structure of C–S–H phases with low C:S ratios consist of non-bridging silica tetrahedra connected by bridging silica tetrahedra (Fig. 2a). The Si in each non-bridging silicate tetrahedron shares all its oxygens with other Si ions or with Ca in the octahedral plane. The Si in the bridging silica tetrahedra, however, have two unshared oxygen atoms neutralised by H+, thus forming a silanediol (≡Si-(OH)2 group). Furthermore, the silicate end-groups, that is silicates connected to only one other silicate group (Si[Q1]), also carry an unshared oxygen. With increasing pH the silandiol groups in the interlayers and the silanol end-groups at the edges become dissociated and form potential sorption sites for Ca2+ or other cations, e.g. actinides and lanthanides (Richardson, 2004).

5 PSI Bericht 18-02

Fig. 2a: Schematic of the tobermorite-based structure of a C–S–H phase with a C:S ratio of

0.66 and a maximum degree of protonation of the silica chains (adapted from Richardson, 2008).

Fig. 2b: Schematic of the tobermorite-based structure of a C–S–H phase with a C:S ratio of

1.5 obtained by removal of the bridging tetrahedra and neutralising the negative charges of the silanol groups by Ca2+ ions (adapted from Richardson, 2008).

In this study three types of uptake mechanisms for actinide and lanthanide ions are envisaged: 1) Binding to the silandiol groups in the interlayers and on planar surfaces, and 2) sorption to silanol end-groups at the surface of C–S–H particles. 3) In addition, actinide and lanthanide ions may become incorporated in C–S–H phases by substitution for Ca2+ in the Ca-O layers. The batch sorption experiments described in this report aim at evaluating the influence of each of these uptake mechanisms on the sorption behaviour of the different actinides and lanthanides.

PSI Bericht 18-02 6

The sorption capacity of the edge and planar silanol groups of C–S–H phases can be estimated based upon the measured surface area after Brunauer, Emmett and Teller (Brunauer et al., 1938) (BET surface area) (~148 m2 g-1Tits et al., 2006) and accepting a surface site density of 4.8 sites per nm2 calculated by Labbez et al. (2006) on the assumption that the C–S–H structure closely resembles the structure of 14Å tobermorite. This site density is comparable to the density of surface hydroxyl groups on amorphous SiO2 (5 sites per nm2) reported by Schindler (1984). With this site density, a sorption capacity of 1.2 eq kg-1 is derived. Note that the sample treatment applied prior to a BET measurement (2 hours under vacuum at 80°C) can change the C–S–H structure dramatically and thus the BET surface area. The total sorption capacity of C–S–H phases (i.e. the total number of silanol and silandiol groups on planar, interlayer and edge sites) can be derived based upon the C–S–H structure and the mass of CaO and SiO2 used to synthesize the C–S–H phases (mass of dry C–S–H). The molecular weight of C–S–H phases (mW) is operationally defined as follows:

mWCSH = (C:S)·mW

CaO + mWSiO2, (2.1)

with mWCaO and mW

SiO2 as the molecular weights of CaO and SiO2. respectively. This operational definition means that one mol C–S–H phase contains one mol Si atoms. C–S–H phases with C:S ratios of 0.67 and 1.65 thus have mW values of 96 g mol-1 and 152.4 g mol-1, respectively. Assuming that the structure of a C–S–H phase with low C:S ratio (C:S = 0.67) closely resembles the structure of tobermorite, on average two sorption sites (OH- groups on bridging tetrahedra) can be identified per Si-trimer (three Si atoms) or 2 moles sorption sites per 3 moles of C–S–H phase. Multiplying this value with the mW

C–S–H , the cation exchange capacity (CEC) of the C–S–H phases with C:S = 0.67 is estimated to be 6.9 moles kg-1. With increasing C:S ratio, bridging Si tetrahedra are progressively removed and the OH- groups connected to the bridging tetrahedra are replaced by two new OH- groups on non-bridging tetrahedra next to the position of the missing bridging tetrahedron. Thus, C–S–H phases with a C:S ratio above 1.5 contain on average two sorption sites per Si dimer (two Si atoms). Multiplication with mW

CSH results in a CEC value of ~6.2 moles kg-1. Both values are approximately a factor 5 higher than the CEC values on the planar and edge sites obtained with the first approach.

The number of Ca2+ sites available for substitution in the Ca-O sheets can be estimated in a similar way: Assuming again close resemblance between the tobermorite structure and the C–S–H structure, a C–S–H phase with a C:S ratio of 0.67 (all bridging tetrahedra in the silica chains are present) contains two Ca2+ ions in the Ca-O layer per Si-trimer or 2 moles Ca2+ per 3 moles C–S–H . Above a C:S ratio of 1.07, (all bridging tetrahedra are removed), C–S–H phases contain equal amounts of Ca in the Ca-O layers and Si in the silica chains or 2 moles Ca2+ per 2 moles C–S–H . Multiplication with the operational molecular weight of the C–S–H phases leads to values between 6.5 moles kg-1 and 6.9 moles kg-1 Ca that are available for substitution in the Ca-O layer.

7 PSI Bericht 18-02

2.2 Actinide chemistry

The actinides form, together with the lanthanides, the F-elements in the periodic table of the elements. These are elements with 4f subshell (lanthanides) or the 5f subshell (actinides) filled with electrons. Actinides and lanthanides can occur in different oxidation states under repository conditions resulting in positively charged cations with charges varying between +1 and +4. Actinide/lanthanide adsorption mainly takes place by electrostatic forces (ionic bonding) at negatively charged hydroxyl groups of the silanol / silandiol sites on the planar surfaces and edges of C–S–H phases and in their interlayers. As a consequence, the actinide adsorption behaviour strongly depends on the charge, and thus on the redox state of the actinide ions.

Lanthanides have only three different oxidation states (+II, +III, +IV) with +III as the most stable oxidation state in the alkaline conditions prevailing in a cement-based repository (10 < pH < 13.3). Unlike the lanthanides, the actinides show a great variety in oxidation states from +III to +VI under alkaline conditions. +III is the most stable oxidation state for the trans-Americium elements (e.g. Cm) but also Pu exists in the oxidation state +III under strongly reducing conditions. Because of their similar chemistry, lanthanides are often used as chemical analogues in research on trivalent actinides. In contrast to the latter, most lanthanides have stable isotopes and are thus much easier to handle in the laboratory. It is for this reason that this report contains a large section dealing with the sorption of Eu(III) onto cementitious materials. +IV is the main oxidation state of Thorium. This actinide does not play an important role in the inventories of radioactive waste repositories but it is commonly used in actinide research as a chemical analogue for Uranium, Neptunium and Plutonium in their tetravalent states because of its redox stability. Note, however, that the larger ionic radius of Th(IV) (rTh(IV) = 1.05 Å, CN = 8) compared to the ionic radii of the other tetravalent actinides (rU(IV) = 1.00 Å, CN = 8; rNp(IV) = 0.98 Å, CN = 8; rPu(IV) = 0.96 Å, CN = 8) (Shannon, 1976) results in much weaker hydrolysis as illustrated by e.g. the first hydrolysis constants of Th(IV) (log *β°(1,1) = -2.5) compared to U(IV) (log *β°(1,1) = -0.54), Np(IV) (log *β°(1,1) = 0.55) and Pu(IV) (log *β°(1,1) = 0.6) (Thoenen et al., 2014). In the same way, the larger rTh(IV) may result in a weaker tendency for Th(IV) to form surface complexes with the –OH groups of silanol- and silandiol groups on the surface of the C–S–H phases. The use of Th(IV) as a chemical analogue for the other tetravalent actinides in this study may therefore result in an underestimation of the Rd values for the sorption of this category of actinides onto C–S–H phases.

In alkaline conditions, Uranium, Neptunium and Plutonium can occur in the oxidation states +III (only Pu), +IV, +V and +VI. In their higher oxidation states (+V, +VI), these actinides form cationic actinyl species ( n

2AnO + , n=1,2) whereby two oxygens form strong covalent bonds with the central actinide (linear O=An=O) (Kaltsoyannis, 1999; Edelstein et al., 2008).

PSI Bericht 18-02 8

2.3 State of the art on actinide sorption on cement minerals

Actinides are classified as hard Lewis acids (Pearson, 1963) and thus, they exhibit a strong affinity for oxygen-containing hard bases, such as the surface hydroxyl groups of metal oxides and C–S–H phases (e.g. Choppin, 1983). These surface hydroxyl groups have oxygen atoms with lone electron pairs, which can act as electron donors in coordination reactions resulting in ionic bonding. Thus, actinide sorption onto C–S–H phases, and more generally hardened cement paste (HCP), is expected to be strong. Neglecting steric effects, the bond strength in ionic bonds correlates roughly with the effective charge of the metal cation (Choppin, 1983, 1984). Hence actinide sorption is expected to decrease with decreasing effective charge of the actinides (An) in the order:

An 4+(4) > AnO22+(3.3) > An3+(3) > AnO2

+(2.2) (2.2)

Numbers in brackets represent the effective charge of the respective actinides (Choppin, 1983, 1984). The “effective charge theory” proposed by Choppin (1983, 1984) does not take into account the influence of the aqueous speciation on actinide sorption. Only few studies have been devoted to the influence of hydrolysis on the sorption behaviour of actinides under alkaline conditions. Yamaguchi et al. (2004) observed that distribution coefficients of U(VI), Np(V) and Sn(IV) in amorphous SiO2 suspensions and in γ-Al2O3 suspensions decreased with increasing pH in the range 9 < pH < 13.5. Tits et al. (2014b) observed a similar sorption behaviour for Np(V) and Np(VI) on titanium dioxide in the pH range between 10 and 14. In cementitious environments, as well, there are indications for an effect of progressing hydrolysis on the sorption of pentavalent and hexavalent actinides. Ochs et al. (2016) compiled literature sorption data for Np(V) and U(VI) on different types of cements in different stages of cement degradation and on C–S–H phases in the pH region 10 < pH < 13.5. The Rd values reported in their review show a clear tendency to decrease with increasing pH value, which was explained as an effect of competition with Ca for the sorbing sites. Indeed, in the pH region 10 < pH < 12.5,the Ca concentration increases with increasing pH due to the higher Ca solubility of the C–S–H phases. Above this pH, however, the Ca concentration decreases again due to decreasing Ca solubility of portlandite in cement. Thus, in the pH region 12.5 < pH < 13.3, competition with Ca cannot explain the decrease in Rd values with increasing pH as noted by Ochs et al. (2016). An alternative explanation could be the lower affinity of more hydrolysed actinide species for sorbing sites on the surfaces of cement phases. Tits et al. (2014b) applied the concept of electrostatic inter-ligand repulsion proposed by Neck and co-workers (Neck and Kim, 2000, 2001; Fanghänel and Neck, 2002) to explain the decreasing affinity of pentavalent and hexavalent actinides for surface sorption sites. The concept predicts that, for each oxidation state of a given actinide, only a limiting number of OH groups, nlimit, can fit in its first coordination sphere. Neck et al. (2000) estimated nlimit values for the trivalent, tetravalent, pentavalent and hexavalent actinides, to be 4, >4, 2 and 4, respectively. The formation of a surface complex results in an additional bond between a surface oxygen and the metal center, thus increasing the total number of ligands in the first coordination sphere (-OH groups + surface oxygen) by one ligand. This means that only hydrolysed actinides with a hydrolysis number lower than nlimit can bind to a sorption site.

Studies on the retention of actinides by cementitious materials in the past have focused predominantly on surface complexation as the relevant uptake process (e.g. Evans, 2008 and references therein). Surface complexation reactions on C–S–H surfaces mainly involve coordination with negatively charged silanol groups on planar sites (sites on interlayer cleavage faces) and at the edges of the C–S–H phases. However, in the past years, evidence has been provided that other immobilization processes, such as incorporation into the solid matrix, probably driven by the high recrystallization rates of C–S–H phases, may play an important role

9 PSI Bericht 18-02

in actinide retardation. With the help of spectroscopic techniques, such as luminescence spectroscopy and X-ray absorption spectroscopy (XAS), several research groups have provided ample evidence that actinides sorb onto C–S–H phases via a two-step process: a rapid surface complexation process is followed by a slow incorporation into the C–S–H matrix (Pointeau et al., 2001; 2003; Schlegel et al., 2004; Stumpf et al., 2004; Mandaliev et al., 2010a; 2010b; Gaona et al., 2011; Tits et al., 2011; 2013a; 2015). Incorporated actinide ions were located both in the C–S–H interlayers (actinides in various oxidation states) and in the Ca-O layers (trivalent actinides).

Trivalent actinides and lanthanides such as Cm(III), Eu(III) and Nd(III) were found to sorb primarily via a rapid surface complexation process, followed by slow incorporation into the C–S–H interlayers and Ca-O layers (Tits et al., 2003; Schlegel et al., 2004; Stumpf et al., 2004; Mandaliev et al., 2010a; 2010b; Macé et al., 2013).

The large number of neighbouring Si atoms (3 – 6) and Ca atoms (8 – 12) observed in XAS investigations indicate that tetravalent Np was predominantly incorporated in the structure of C–S–H phases (Gaona et al., 2011). The coordination environment of this incorporated species appears to depend on the C:S ratio of the C–S–H phases: Increasing C:S ratios results in sorbed Np(IV) species with lower Si coordination numbers and higher Ca coordination numbers. This led the authors to the conclusion that Np(IV) is bound in the C–S–H interlayers. In another paper, Gaona et al. (2013b) observed that the Np(V)-Oax distances of NpVO2

+ sorbed on C–S–H phases were significantly longer than the axial oxygen distances in the NpVO2

+ aquo ion. This indicates that all H2O molecules were removed from the NpVO2

+ coordination sphere upon sorption. The authors concluded that this is an indication that also Np(V) is incorporated in the C–S–H matrix.

In the case of the hexavalent actinides a two-step process similar to that observed for the trivalent actinides was reported: Tits and co-workers applied luminescence spectroscopy to discern the uptake process of U(VI) by C–S–H phases. They identified two sorbed U(VI) species on U(VI)-loaded C–S–H samples (Tits et al., 2011; 2015): A surface complex having spectroscopic characteristics similar to U(VI) surface complexes on rutile, and an incorporated species characterized by a strongly red-shifted luminescence spectrum. Upon increasing reaction time, the spectrum of the incorporated species became progressively dominant. XAS studies could not confirm this observation because backscattering contributions from Ca and Si neighbours in the XAS spectra were too weak to allow unambiguous identification of the coordination sphere.

Several mechanisms may be responsible for this incorporation process such as recrystallization via Ostwald ripening or solid state diffusion (Curti et al., 2010). Ostwald ripening describes a process in aqueous solution by which smaller particles preferentially dissolve providing material for the growth of larger particles (Steefel and Van Capellen, 1990; Thanh et al., 2014). The driving force for this process is the decrease in surface free energy with increasing particles size. Solid state diffusion is a process by which ionic species migrate through a solid phase driven by a concentration gradient. In common solid phases, solid state diffusion processes are generally too slow (diffusion coefficients are about 10 orders of magnitude lower than in aqueous solution) to explain incorporation processes (Curti et al., 2010).

2.4 Description of the experimental studies

The present report contains a synthesis of all the wet chemistry studies on actinide and lanthanide sorption onto cementitious materials carried out at PSI. The wet chemistry studies

PSI Bericht 18-02 10

include 1) solubility tests of actinides in various types of cement pore waters, 2) recrystallization experiments with C–S–H phases, 3) sorption experiments, 4) co-precipitation experiments and 5) desorption experiments.

1) Actinide solubility tests in cement pore waters

Thermodynamic modelling using thermodynamic databases (TDB’s) allows the speciation and solubility of actinides in various cement pore waters to be predicted. The accuracy of these predictions depends on the quality of the TDB’s applied. Note that the thermodynamic constants listed in these TDB’s are still incomplete and thermodynamic constants for many actinide complexes and solids potentially relevant in highly alkaline cementitious environments (e.g. ternary Ca-OH-An complexes, An-silicate complexes) are still missing. For a correct interpretation of sorption experiments a detailed knowledge of the actinide speciation and the actinide solubility under the chemical conditions of the sorption experiments is crucial. Otherwise, e.g. a decrease in aqueous actinide concentration due to precipitation could be falsely interpreted as sorption on a cement phase. To avoid such misinterpretations, solubility tests were carried out to determine experimental solubility limits in the cement pore waters planned to be used in later sorption studies.

2) Recrystallization experiments with C–S–H phases

Spectroscopic studies carried out in-house and reported in the literature have shown that most actinides become incorporated in the matrix of cement minerals with time during sorption experiments. Potential mechanisms resulting in actinide incorporation are matrix diffusion or co-precipitation during recrystallization processes. The former mechanism is known to be much too slow to explain the actinide/lanthanide uptake observed in the sorption experiments. The kinetics of recrystallization of C–S–H phases can be followed with the help of 45Ca uptake experiments. In these experiments, a 45Ca tracer is added to a suspension containing a C–S–H phase in equilibrium with its C–S–H solution, and the uptake of the 45Ca tracer is followed with time. As the C–S–H phase is in equilibrium with its solution, 45Ca uptake can only take place by recrystallization (Ostwald ripening). The 45Ca uptake with time thus is a measure of the C–S–H recrystallization kinetics in the experiment. The C–S–H recrystallization kinetics will be compared with the kinetics of the sorption of various actinides onto C–S–H phases to explore whether a correlation exists between the two processes.

3) Sorption experiments

The term “sorption” is defined in this report as any process by which an aqueous species is temporarily or permanently immobilized by a solid phase; i.e. “sorption” includes processes such as “surface complexation”, “ion-exchange”, “surface precipitation” and “incorporation”. A sufficiently detailed understanding of the sorption behaviour of the actinides is compulsory for a reliable prediction of the mobility of actinides in the cementitious near-field of a L/ILW repository. “Sorption behaviour” includes the effects of parameters such as reaction time, actinide concentration, pH, pore water composition, and redox conditions, but also the reversibility of the sorption reaction. The sorption studies described below were designed to gain insight in the effect of each of these parameters on actinide sorption.

C–S–H phases are characterized by very high reactive surface areas providing large numbers of surface sites accessible to actinide cations for surface complexation or cation exchange. In addition, it will be shown in chapter 4.3. that C–S–H phases are very reactive materials with high recrystallization rates. These characteristics make C–S–H phases ideal materials for both surface complexation and incorporation processes. To determine whether an actinide sorbs on

11 PSI Bericht 18-02

C–S–H phases by surface adsorption rather than incorporation, its sorption behaviour on C–S–H phases is compared to that on Titanium dioxide (TiO2). In contrast to cement phases, TiO2 is not reactive in alkaline environments and therefore surface complexation with the hydroxyl groups on the TiO2 surfaces is the only mechanism enabling actinide binding on this material. Note that TiO2 exhibits significant photocatalytic properties and thus may influence the redox state of redox-sensitive actinides. For example few studies have reported that TiO2 is capable of reducing U(VI) to U(IV) under UV irradiation (Eliet and Bidoglio, 1998; Bonato et al., 2008; Odoh et al., 2012) and that limited photocatalytic reduction of U(VI) may occur under daylight conditions (Kim et al., 2015). Such photocatalytic effects, however, only occur under very specific chemical conditions (strong UV irradiation, presence of electron donors,…) not encountered in the present study. Hence it is assumed that photocatalytic effects can be neglected in the present study.

Sorption kinetic studies investigate the time needed for a sorption reaction to reach equilibrium. Short reaction times are an indication for rapid sorption process such as ion-exchange or surface complexation whereas longer reaction times are an indication that slower processes such as incorporation may control sorption.

The effect of the actinide equilibrium concentration in solution is studied with the help of sorption isotherms and provides information about the linearity of the sorption process. Linear sorption means that the amount of actinide species sorbed on the solid is proportional to the equilibrium solution concentration of the sorbing actinide species and is an indication that one single actinide species sorbs on one single type of sorption sites. Non-linear sorption indicates that either more than one actinide species or more than one sorption site are involved.

The effect of pH and pore water composition provides information on the effect of actinide speciation (hydrolysis) on the sorption process; i.e. the type of sorbing species and competition between the sorption process and hydrolysis in solution.

The redox conditions in the experiments determine the redox state and thus the charge of the actinide. The effect of the charge of actinide ions on sorption was described in chapter 4. The redox conditions in the sorption experiments were fixed by using redox buffers. Na-hypochlorite (NaClO) was used to produce oxidizing conditions:

2ClO H O 2e Cl 2OH− − − −+ + → + E° = 0.89 V (2.3)

Na-dithionite (Na2S2O4) was used to produce reducing conditions:

2 23 2 42SO 2e S O 4OH− − − −+ ⇔ + E° = -1.13 V (2.4)

Gaona et al. (2013a) and Rojo et al. (2013) showed that Na-dithionite and Na-hypochlorite have no influence on the sorption behaviour of tetravalent and hexavalent actinides. X-ray absorption near edge spectroscopy (XANES) spectra of C–S–H pastes containing much higher 237Np concentrations (>10-5 M) in the pH range 9 < pH < 13.5 confirmed that these reducing/oxidizing agents were able to keep Np in the tetravalent, respectively the hexavalent states (Gaona et al., 2011; Gaona et al., 2012b). Thus, it may be safely assumed that Na-dithionite and Na-hypochlorite are capable of controlling the redox state of 239Np at the much lower concentrations used in the present experiments.


4) Co-precipitation experiments

Co-precipitation experiments have been carried out to simulate the incorporation of actinides in C–S–H phases. The set-up of such experiments provides better conditions for the actinide to become incorporated in the matrix of the resulting solid phase. Indeed, in co-precipitation experiments, the actinide under investigation is added at the beginning of the synthesis of the C–S–H phases and, thus, it can easily occupy the thermodynamically most favourable position in the crystal matrix. On the other hand it is known, however, that fast co-precipitation from strongly oversaturated solutions may result in entrapment of the actinide in a less thermodynamically stable position (Heberling et al., 2014).

5) Desorption experiments:

Immobilization of aqueous metal species by a solid phase may occur via several sequential processes, e.g. rapid surface complexation followed by surface migration from weak sorption sites to strong sorption sites, or also, followed by a slower incorporation process, e.g. via recrystallization. Macroscopic indications for such sequences of rapid and slow processes can be found with desorption experiments. Differences in sorption and desorption kinetics (sorption hysteresis) are an indication that a rapid surface complexation is followed by a slower uptake process. Such sorption hysteresis effects are, however, not an indication for an “irreversible” process. Taking into account the extremely long timescales dealt with in nuclear waste disposal, all slow incorporation processes are considered to be thermodynamically reversible processes as well. During incorporation processes a new solid phase is formed (e.g. a solid solution) defined by an equilibrium constant (e.g. solubility product). The combined surface sorption and incorporation kinetics during the sorption test may be different from the combined surface desorption and recrystallization kinetics during the desorption test resulting in sorption hysteresis.


3 Materials and methods

3.1 Materials

Merck “Pro analysis” chemicals and deionized, decarbonated water (MilliQ water) generated by a MilliQ Gradient A10 System (Millipore, Bedford, USA) were used for all the experiments reported in this work. The MilliQ water was de-aerated by purging thoroughly with N2 for at least 2 hours followed by equilibration with the inert N2 glove box atmosphere for at least one week before use to minimize CO2 (and carbonate) impurities. The centrifuge tubes used for the wet chemistry experiments were washed, left overnight in a solution of 0.1 M HCl, and thoroughly rinsed with MilliQ water. All experiments were carried out in glove boxes under N2 atmosphere (pO2, pCO2 < 2ppm).

3.1.1 Experimental cement pore waters Sorption studies with HCP were performed in artificial cement pore water solutions (ACWHCP-I). The composition of ACWHCP-I is based on an estimate of the pore water composition in HCP before any degradation of the material had occurred (Wieland et al., 1998; 2006). A detailed description of this ACWHCP-I and its preparation is reported elsewhere (Wieland et al., 2006). Briefly, an alkaline solution containing 0.114 M NaOH and 0.18 M KOH was prepared. An excess of Ca(OH)2 and CaCO3 was added and this suspension was shaken for at least 1 week. Subsequently, the suspension was filtered through a 0.1 μm Criticap filter (Gelman Science) under CO2-free conditions to remove all particulate material. Then 0.426 g L−1 Na2SO4 and 8.55·10−3 g L−1 Al2(SO4)3 was added to the filtrate to obtain an Al concentration of 5·10−5 M and a SO4

2− concentration of 3·10−3 M. Stability tests showed that the composition of this ACWHCP-I remained unchanged in contact with HCP over a period of at least 4 weeks (Wieland et al., 1998).

Aged ACWHCP-II was prepared by immersing crushed HCP in MilliQ water at a solid-to-liquid (S:L) ratio of 1:40 which caused the cement paste to be degraded to the second stage of the cement degradation (Ca(OH)2 saturated solution). Ca is the main element in this stage of cement degradation, which is controlled by portlandite solubility, while millimolar concentrations and below were achieved for the other elements (Na, K, Al, S).

Sorption studies with C–S–H phases were performed either in a simplified artificial cement pore water (ACW) or in alkali-free C–S–H solutions. ACW is a simplified ACWHCP-I solution: pH and Na and K concentrations are identical but the Ca and Si concentrations of the pore water are fixed by the solubilities of the respective C–S–H phases. The Al3+ and SO4

2- concentrations in this pore water are set to zero. The alkali-free C–S–H suspensions solely contain Ca and Si and the concentrations of these cations as well as the pH are controlled by the solubilities of the respective C–S–H phases.

3.1.2 Titanium dioxide Titanium dioxide (TiO2) (Aeroxide P25) was obtained from Evonik Industries AG (Germany). This material is a mixture of 86% anatase and 14% rutile and has a specific surface area of 56±1 m2 g−1 (Schmidt and Vogelsberger, 2009; Comarmond et al., 2011). TiO2 suspensions were prepared by mixing the appropriate amounts of Aeroxide P25 with 0.1M NaCl solutions to obtain the S:L ratios required for the batch sorption tests. The pH of the suspensions was adjusted with NaOH solutions having concentrations of 0.01 M, 0.3 M and 1.0 M. The pH of


the suspensions had to be adjusted several times over a period of 3 days to compensate for deprotonation reactions on the TiO2 surface. TiO2 suspensions with pH values above 13 were prepared in pure NaOH without 0.1M NaCl as background electrolyte.

3.1.3 C–S–H phases C–S–H phases with varying C:S ratios (0.65 < C:S < 1.65) were synthesized in MilliQ water (alkali-free C–S–H phases) and in ACW following a procedure described elsewhere (Atkins et al., 1991; Tits et al., 2006). Briefly, AEROSIL 300 (SiO2) (Evonik Industries AG, Germany) was mixed with CaO in polyethylene bottles to give target C:S ratios between 0.6 and 1.82. To this, ACW or Milli Q water was added to achieve S:L ratios between 5·10−3 kg L−1 and 2·10−2 kg L−1 (based upon the weight of CaO and SiO2 used for the synthesis). Note that in this report, the mass of C–S–H phase refers to the sum of the masses of CaO and SiO2 (kg L-1) used for the synthesis of the respective C–S–H phase:

mC–S–H = mCaO + mSiO2 (3.1)

The suspensions were mixed and homogenized with a shear mixer. After equilibration for at least two weeks on an end-over-end shaker, the C–S–H suspensions were ready for further use in sorption experiments (fresh C–S–H suspensions). Preliminary kinetic tests showed that the solution compositions of the C–S–H suspensions did not change after two weeks equilibration (see section 7.1). Two 40 mL aliquots of the suspensions were sampled, centrifuged for 15 min at 10,000 g (max) and the supernatant solutions analysed for Ca, Si, Na, K concentrations with an inductively coupled plasma optical emission spectrometer (ICP-OES). The composition of the solution was further checked in some of the sorption experiments. The C–S–H suspensions were filtered through Whatman ashless 541-grade filter paper using a Buchner funnel with vacuum pump. The filtrates were collected and stored for analysis of their ionic composition. The filter cakes were washed with MilliQ water to remove the remaining salts and dried to constant weight in a desiccator over a saturated CaCl2 solution which maintains a relative humidity of 30%. The contact time of the slurries with MilliQ water was kept as short as possible (i.e.i.e. less than 1 min to minimize changes in composition).

In some sorption experiments S:L ratios lower than 5·10-3 kg L-1 were required. In these cases, an aliquot of the C–S–H suspension was centrifuged on a L7-35 Ultracentrifuge (Beckmann Instruments, Inc.) (15 min at 10,000 g (max)) and the supernatant solution was used to dilute the remaining C–S–H suspension to the appropriate S:L ratio.

In suspensions with low S:L ratios, the final C:S ratios of the synthesized C–S–H phases could not be directly deduced from the stoichiometric ratios of CaO and SiO2 used to synthesize the C–S–H phases because small portions of CaO and SiO2 dissolved to establish the respective Ca and Si contents of the C–S–H equilibrium solutions. The actual C:S ratio was then calculated from the starting CaO and SiO2 concentrations and the measured Ca and Si concentrations in the equilibrium C–S–H solutions.

3.1.4 Hardened cement paste The cement samples were prepared from a commercial sulfate-resisting cement (Type CPA 55 HTS; Lafarge, France), denoted as HTS (haute teneur en silice). A special manufacturing procedure was developed to produce fully hydrated HCP under CO2-free conditions (Wieland et al., 2006). For the batch sorption experiments, a powder material was prepared by crushing the


bulk material in a mortar under CO2-free conditions and sieving the material to collect the fraction <70 μm. The BET surface area of the cement powder was determined to be 46±4 m2 g-1.

3.2 Analytical methods

3.2.1 Radiotracers and radio-assay Radiotracers were used to measure ion distribution ratios in the batch sorption experiments and the C–S–H recrystallization experiments. 45Ca, 152Eu, 228Th and 233U tracers were obtained from Imatom GmbH, Switzerland. A 237Np stock solution was kindly provided by M. Altmaier (KIT-INE, Karlsruhe, Germany). 239Np tracer solutions were prepared according to a procedure described by Sill (1966). The separation of this short-lived actinide isotope from its parent 243Am (Eckert & Ziegler Isotope Products, USA) was achieved by extraction into long-chain amines (tri-iso-octylamine in xylene) followed by back-extraction with 0.01 M HCl.

152Eu, 239Np, 237Np, 233U, 228Th and 45Ca activities in diluted suspensions and in solution were determined by liquid scintillation counting using a Perkin Elmer Tri-CarbTM A2750 CA liquid scintillation analyser (LSA) equipped with alpha-beta discrimination option. Prior to liquid scintillation analysis, 5 mL sample aliquots containing 152Eu, 239Np or 45Ca were mixed with 15 mL Ultima Gold XR scintillator (Perkin Elmer Inc., USA) and 3 mL samples containing 237Np, 233U, 228Th were mixed with 17 mL Ultima Gold AB scintillator (Perkin Elmer Inc., USA). 228Th samples were stored for at least one month prior to analysis to allow secular equilibrium of 228Th with its daughters to be established. This secular equilibrium was disturbed after uptake of the tracer by the solid due to the much higher sorption of 228Th (Rd = ~ 105 L kg-1) compared to its daughter 224Ra (102 < Rd < 104 L kg-1; Tits et al., 2006). Activities in concentrated suspensions were not determined with liquid scintillation because there is a risk for artefacts due to radiation quenching effects caused by the particles in the suspensions. Instead, suspensions were analysed by g-counting using a Packard Cobra-II model 5003 gamma counter equipped with a 3 inch NaI through-hole crystal detector.

3.2.2 Chemical analyses Solution compositions were determined using an Applied Research Laboratory ARL 3410D inductively coupled plasma optical emission spectrometer (ICP-OES). 5 mL aliquots of supernatant solutions were diluted with 5 mL MilliQ and acidified with 0.1 µL concentrated HNO3.

The pH was determined using a combination glass pH electrode (Metrohm, Switzerland) calibrated against standard pH buffers (pH = 7.0 - 11.0, Metrohm, Switzerland). The measured pH value (pHexp) was converted to molar H+ concentrations, [H+] (with pHc = -log[H+]) as follows: In solutions of ionic strength I < 0.1 M, the measured pHexp value is related to pHc by

Hc exp HH

apH log[H ] log pH log

++

+

+ = − = − = + g

g , (3.2)

where Ha + and H+g are the H+ activity and the activity coefficient for H+. H+g is calculated with the help of the Davies equation (b = 0.3) (Davies, 1962). In salt solutions of ionic strength I ≥ 0.1 mol L–1, the measured pH value (pHexp) is an operational apparent value related to [H+] by


pHc = pHexp + Ac, (3.3)

where Ac is given as a function of NaCl and CaCl2 concentrations (Altmaier et al., 2003; Gaona et al., 2013a).

The redox potential of the samples was determined using a Metrohm combined redox electrode consisting of a Pt-ring indicator electrode and a Ag/AgCl reference electrode. The performance of the combined redox electrode was regularly checked using a Metrohm redox buffer at +250 mV (vs. Ag/AgCl) at T = 20°C. Voltage readings were corrected to standard hydrogen electrode (SHE) values.

3.2.3 Experimental uncertainties and reproducibility Sorption measurements are subject to several sources of uncertainties giving rise to an overall uncertainty on the sorption values. In the present study the uncertainty on the activity measurements for specific experimental conditions was evaluated following an approach described by Tits et al. (2002). This approach is also valid for the actinide/lanthanide sorption experiments onto C–S–H phases and HCP in the present report. It consists of the definition of a minimum measurable activity in solution (Al,min) and a maximum distribution ratio, Rd,max which are measurable with sufficient precision (relative uncertainty < 10%) in a specific sorption experiment. Al,min and Rd,max depend on the specific conditions of each experiment: Al,min depends on the background radioactivity for each specific radioanalysis. The Rd,max value depends on Al,min and on the S:L ratio used in each specific sorption experiment.

Once the Rd,max value is known for a specific sorption experiment, then, in a second step, the uncertainty (relative 95% confidence interval) on the measured Rd values that fall below the limits of Rd,max, is required. The relative 95% confidence intervals on the equilibrium metal concentrations and on Rd values for typical Eu(III) and Th(IV) sorption experiments on calcite under hyperalkaline conditions were found to be approximately 40% and 67% (Tits et al., 2002). These high uncertainties are mainly related to the very strong uptake, which results in very high ratios (>1000) of the sorbed activity to the activity in the aqueous phase. Incomplete phase separation or resuspension of very small amounts of solid phase during sampling of the supernatant may then give rise to significant overestimation of the activities in the aqueous phase.

In most experiments described in the present report, the number of replicates was too small to allow for a correct estimate of the 95% confidence interval on the Rd value. Although the sorption experiments described in Tits et al. (2002) were carried out on calcite, the experimental set-up and conditions are very similar. It was therefore assumed that the experimental uncertainties on equilibrium activity measurements and on Rd measurements in this report are similar to the uncertainties found by Tits et al. (2002). Therefore, 95% confidence intervals of 40% and 67% were applied on equilibrium concentrations and Rd values, respectively.

3.3 Experiments

3.3.1 Characterization of the C–S–H phases and HCP material C–S–H suspensions prepared as described above were filtered through Whatman ashless 541-grade filter paper using a Buchner funnel with a vacuum pump. The Na, K, Ca, and Si concentrations in the filtrates were determined using ICP-OES. The elemental composition of


the synthesized C–S–H phases was obtained from simple mass balance calculations by taking into account the amounts of CaO and SiO2 added during synthesis and the Ca and Si concentrations measured in the equilibrium solutions after synthesis as previously mentioned (Section 3.1.3.).

The specific surface area of some of the C–S–H phases was determined using the N2-BET method (e.g. Brunauer et al., 1938; Kantro et al., 1967) by measuring multipoint N2 sorption isotherms with a Micrometrics Gemini 2360 surface area analyzer. Note that the specific surface area determined with this method only gives an approximate value of the actual in-situ surface area of the C–S–H phases due to severe treatment of the samples (drying at 80°C under vacuum) prior to the measurement.

Quantitative X-ray diffraction (XRD) measurements of the powders were performed between 1.5° and 96.99° 2θ using a Bragg-Brentano diffractometer (Bruker AXS D8, Cu Kα radiation, automatic divergence slit, graphite monochromator). The step width was 0.03° 2θ and the step counting time was 4 s. The portlandite content was determined from XRD pattern fitting (Müller, 2005). Prior to each XRD measurement a known amount of an internal standard (CaCO3, Fluka) was added to the C–S–H powder material. This mixture was gently homogenized by grinding. Only the XRD lines of the internal standard and the portlandite were fitted with the program BGMN® (Bergmann et al., 1998). The resulting portlandite contents were then recalculated with the known content of the internal standard.

Characterization of the HCP material used in the present study is reported in detail elsewhere (Lothenbach and Wieland, 2006). HCP prepared from the HTS cement at a water to cement ratio of 0.4 consists of mainly C–S–H phases (~44 weight (wt)%), portlandite (~19 wt%), ettringite (~9 wt%), AFm phases (~7 wt%) with some minor phases such as hydrotalcite, calcium carbonate, and non-reacted clinker minerals.

3.3.2 Recrystallization experiments with C–S–H phases The recrystallization of the synthetic C–S–H phases was studied with the help of isotopic exchange experiments. In these experiments, a radioactive isotope of one of the major ions in the C–S–H phases (45Ca) was added to the C–S–H suspensions and the uptake of this radiotracer by the C–S–H phases was determined as a function of time. The rationale behind isotope exchange experiments is that this isotope becomes progressively incorporated into the solid phase during the recrystallization process and further, that the rate of incorporation is similar to the C–S–H recrystallization rate. The isotope exchange technique and its application to the characterization of recrystallization processes is described in detail elsewhere (Curti et al., (2010). The goal of the isotopic exchange experiments was to provide evidence that the actinide uptake by C–S–H phases could be driven by recrystallization processes. Note that no re-crystallization experiments were carried out with HCP as several Ca containing cement phases exist in the complex cement matrix, i.e. portlandite, C–S–H , ettringite, AFm phases, calcium carbonate etc., which limits an unambiguous interpretation of recrystallization data.

3.3.3 Preliminary solubility tests of actinide/lanthanide solutions A sufficiently detailed knowledge of the aqueous speciation is a prerequisite to allow detailed interpretation of sorption experiments with actinides and lanthanides under highly alkaline conditions. For example, the solubility of the actinides/lanthanides under investigation must be known to allow sorption experiments to be designed in such a way that artefacts caused by precipitation are avoided. Under highly alkaline conditions, the detection limit of some


actinides/lanthanides is very close to the solubility limit. In addition, thermodynamic solubility data are always subject to uncertainty and stability constants of some potential complexes are still unknown (e.g. ternary Ca-An-OH and Ca-Ln-OH complexes and ternary An-OH-silicate and Ln-OH-silicate complexes). Finally, the apparent “solubility” observed in an experiment does not necessarily correspond to the solubility calculated based on thermodynamic data. Artefacts caused by short-time local supersaturation during addition of the actinide/lanthanide tracer to a C–S–H suspension may result in the formation and precipitation of metastable colloids which only very slowly dissolve again at a later stage. Therefore, the apparent solubility of actinide/lanthanide solutions was checked under the chemical conditions (pH, solution composition) established in the sorption experiments.

Solutions containing increasing total tracer concentrations (concentration range depends on the actinide/lanthanide under investigation) were prepared. One series of samples was shaken for 1 day, the others for a longer period of time (30 days or 60 days, respectively). After the ageing periods, the radionuclide activities in solution were measured before and after centrifugation (1 h at 28,000 rpm (95,000 gmax)). Applying this stepwise procedure allowed the radionuclide fraction in "true" solution to be distinguished from the radionuclide fraction present as colloidal material.

Application of Eq.( 3.4) allows estimating the minimum size of the particles that are settled at a particular centrifugation speed. Eq. (3.4) is derived from Stokes Law by replacing the gravitational acceleration, g, by w2r, where w is the angular velocity of the particles in rad per second and r is the radius of the particle (Thorstenfelt et al., 1988):

( )

2

122

p w

R9 lnRr

2 t

⋅ η⋅ =

⋅ r − r ⋅ w ⋅ ∆. (3.4)

∆t: centrifugation time = t2 – t1 (s) R1: radial position at time t1, the start of the centrifugation R2: radial position at time t2, the end of the centrifugation r: radius of the particle (cm) w: angular velocity (w = 2·π·rpm/60, rad s-1), where rpm = revolutions per minute rp: particle density (g cm-3) rw: density of water (g cm-3) η: dynamic viscosity (mPa·s or g s-1 cm-1)

In the particular case of actinide/lanthanide colloids in aqueous alkaline solutions at 25°C, we assume that η = 0.89·10-2 g s-1 cm-1. The density of water is set to 1.0 g cm-3 and the density of the actinide colloids, rp, set to a value of 9 g cm-3. This value was determined by Rundberg et al (1988) for PuO2 colloids in a solution at pH = 2 - 3. The density of the actinide/lanthanide colloids in our experiments may vary significantly depending on the type of actinide/lanthanide and the stoichiometry. The value of 9 g cm-3 is only a very rough approximation. A centrifugation speed of 28,000 rpm gives an angular velocity, w, of 2.93·103 rad s-1. The values for R1 (7.0 cm) and R2 (10.8 cm) were derived from the characteristics of the Beckman centrifuge rotor used. Applying this set of values to equation 3.4 shows that any actinide/lanthanide colloids with a diameter greater than ~ 3 nm should settle during the centrifugation step.


3.3.4 Sorption experiments TiO2, C–S–H and HCP suspensions were prepared in 40 mL centrifuge tubes as described earlier by immersing appropriate amounts of the solid material in their equilibrium solution (Section 3.1.3.). Aliquots of actinide/lanthanide stock solutions (in 0.1 M HNO3) were added stepwise to the fresh suspensions while strongly stirring the suspensions to achieve the concentrations required in each experiment. The actinide/lanthanide concentrations of these stock solutions were chosen in such a way that only small volumes (< 1 mL) had to be added to reach the total concentration required in the experiments.

In the Np(IV) and Np(VI) sorption studies, the redox state of the actinide tracer was controlled by addition of 5·10-3 M Na-dithionite (Na2S2O4) or 10-2 M Na-hypochlorite (NaClO), respectively.

Few sorption experiments with C–S–H phases were carried out with material that had been dried prior to use. In these experiments the C–S–H material was re-suspended in a solution having a composition in equilibrium with the suspended C–S–H phase, homogenized with a shear mixer and equilibrated for at least one week prior to addition of the radiotracer.

The centrifuge tubes were shaken end-over-end for an appropriate period of time (varying time periods in the case of the kinetic experiments, two weeks for the other types of experiments). After equilibration, duplicate samples were taken from each homogeneous suspension and the total actinide/lanthanide activity in the suspensions was determined with radioassay. Phase separation was done by centrifugation (1 h at 95,000g). Application of equation 3.4, and assuming the particle density of C–S–H phases to be 2.6 g cm-3 (Allen et al., 2007) shows that this centrifugation step allows solid C–S–H particles down to a diameter of ~6 nm to be separated from solution. After centrifugation, triplicate samples were withdrawn from the supernatant solution and analysed either by ICP-OES or by LSA. Sampling of the supernatant solution of C–S–H suspensions after phase separation by centrifugation turned out to be a challenging task. In many cases, phase separation was incomplete and a thin film of very fine particles was observed on the surface of the supernatant solution. Occasionally, some particles were mixed into the supernatant solution during sampling leading to erratic measured radionuclide activities in the supernatant solutions due to the strong interaction of the actinides / lanthanides with particles.

Sorption samples with 239Np were equilibrated for only three days, because the short 239Np half-life (2.355 days) and the low activities remaining in the supernatant after sorption in many experiments did not allow longer equilibration times. Sorption kinetics tests (Section 3.1) showed that such short equilibration times may result in an underestimation of the Rd value of 50% at maximum. This underestimation is still within the 95% confidence interval applied for all Rd values presented in this study.

3.3.5 Co-precipitation experiments with C–S–H phases Co-precipitation experiments were carried out by adding appropriate volumes of actinide/lanthanide solutions at the beginning of the synthesis procedure used for the C–S–H phases. Equilibration, phase separation and radioassay were similar to the procedures described above for the sorption experiments


3.3.6 Desorption experiments on C–S–H phases Actinide/lanthanide desorption tests onto C–S–H phases synthesized in MilliQ water and in ACW were carried out after sorption of the actinide/lanthanide. Actinide/lanthanide containing C–S–H suspensions were equilibrated for 1 and 30 days using different initial actinide/lanthanide concentrations. After phase separation by centrifugation (1 h at 95,000g), the supernatant solutions were replaced by an actinide/lanthanide-free solution, which had been pre-equilibrated with a C–S–H phase of the same chemical composition (i.e. C:S ratio). The solids were re-suspended in the fresh solution using a shear mixer and the suspensions were then shaken end-over-end over a period of time up to 30 days. Phase separation and radioassay were similar to the procedures described earlier for the sorption experiments. For some desorption tests, up to three sequential desorption steps were carried out and after each replacement of the aqueous phase, the re-suspended materials were allowed to equilibrate for three days.

3.4 Interpretation of the sorption data

3.4.1 Distribution ratios and concentration calculations The uptake of the actinides/lanthanides by TiO2, C–S–H phases and HCP as well as the uptake of 45Ca by C–S–H phases in the recrystallization experiments was quantified in terms of the radionuclide distribution between solid and liquid phase by measuring the activities of the radionuclides in both the suspension and the supernatant solution. The uptake is expressed in terms of a distribution ratio, Rd (L kg-1), which is the ratio of the amount of radionuclide sorbed, Msorb, (mol kg-1) and the radionuclide concentration in solution, [M]l (mol L-1). Rd values were obtained using the following equation:

{ } susp lsorb sorb

dl l l

M (A A )A V VR[M] A m A m

−= = ⋅ = ⋅ (L kg-1), (3.5)

Asorb is the radionuclide activity sorbed on the solid (Bq·L-1), Al is the activity determined in the supernatant solution (Bq L-1), Asusp is the activity determined in suspension (Bq L-1), V is the sample volume (L) and m is the mass of solid phase in suspension (kg). Again, it is noted that in the case of C–S–H phases the sorption values reported in the following sections will be given relative to the “dry weight”, i.e. the weight of the CaO and SiO2 used in the synthesis of the respective C–S–H phase.

In many sorption tests, actinide/lanthanide sorption was so strong that the use of very low S:L ratios (<10-3 kg L-1) was required to obtain activities in the supernatant solutions above the minimum measurable activity Al,min. As a result of these low S:L ratios, part of the actinide/lanthanide activity was sorbed on the wall of the centrifuge tubes during the sorption experiments. To correct for this wall sorption effect, Asusp was used in the calculation of the Rd values.

In some experiments with higher suspension concentrations, β activities could not be measured properly due to quenching caused the solid particles. In this case, theg activity was measured in the suspension (Asusp

g) and the β activity (Alβ) in the supernatant solution. The counting

efficiencies of the two techniques are different. Therefore, Asuspg has to be converted into

Asuspβ prior to calculation of the Rd value using a conversion factor, f(g→β):


( )SAfSA

βg→β =

g, (3.6)

where SAβ and SAg are the specific β activity and the specific g-activity of the actinide or lanthanide solution (cpm/mol), respectively. SAβ and SAg are obtained from β and g measurements of series of standard solutions with known radionuclide concentrations:

,tot totSA A [M]g g= , ,tot totSA A [M]β β= (cpm mol-1), (3.7)

with Ag,tot, Aβ,tot and [M]tot being the total actinide/lanthanide g and β activities (Bq L-1) and the total actinide/lanthanide concentrations in solution.

Knowing the SAβ, of the actinide/lanthanide solution in a sorption experiment, the actinide/lanthanide concentration in the supernatant solution, [M]l can be calculated as follows:

[ ] ll

AM

SA

β

= (M). (3.8)

The amount of radionuclide sorbed on the solid phase in the sorption experiment, Msorb, is obtained by rearranging Eq. 3.5:

{ } [ ]dsorb lM R M= ⋅ (mol kg-1). (3.9)

3.4.2 Modelling the effect of hydrolysis on sorption A reduction factor, Fred, is defined to assess the influence of the competition between the formation of the non-sorbing nth hydroxy species (n = nlimit) in solution, and sorption. Fred corresponds to the ratio of the Rd value determined in the absence (Rd

0) and in the presence (Rd) of the non-sorbing hydroxy species. Development of the equation for Fred was described in detail elsewhere ((Wieland and Van Loon, 2002; Tits et al., 2014b). A brief summary is given below.

The stepwise hydrolysis of actinides can be written in the following general form (the same approach is valid for the lanthanides):

z (n 1)z nn 2 n 1An(OH) H O An(OH) H− +− +

++ ⇔ + , (3.10a)

with the corresponding conditional stepwise complexation constant, * 'n 1K + :

* ' z (n 1) z nn 1 n 1 nK An(OH) H An(OH)− + + −

+ + = . (3.10b)

“z” is the redox state of the actinide, “n” is the hydrolysis number and ”n+1” is nlimit.

Near the pH value where “n+1” approaches nlimit, the definition of Rd in Eq. (3.5) can be rewritten as:

{ } ( )lim it

lim it

z nz nd n nsorb

R An [An(OH) ] [An(OH) ]−−= + . (3.11)


Combining Eq. (3.10b) and Eq. (3.11) gives:

{ } ( )( )lim it

z n * 'd n nsorb

R An An(OH) 1 K / H− + = ⋅ + . (3.12)

Assuming linear sorption, { } z n 0n dsorb

An [An(OH) ] R− = , Eq. (3.12) can be written as:

( )imit

0 * 'd d nR R 1 K / [H ]+= + . (3.13)

From Eq. (3.13), the following expression for the reduction factor, Fred, can be deduced:

limit* '0

ndred

d

KRF 1R [H ]+= = + . (3.14)

Fred is valid for all solids provided that the hydrolysed species lim it

lim it

z nnAn(OH) − does not sorb, the

sorption of the species z nnAn(OH) − is linear and reversible and the sorbent remains stable

throughout the sorption experiment. The Rd values in the pH range 10 < pH < 14 can then be calculated as follows:

0d

dred

RR

F= . (3.15)


4 Characterization of the C–S–H phases

4.1 Pre-equilibration of solid phases

To determine the time needed to achieve chemical equilibrium in fresh C–S–H suspensions, the evolution of the solution composition of three C–S–H suspensions in ACW was studied as a function of time from the start of synthesis over a period up to 130 days. The measured aqueous concentrations displayed in Fig. 3 show clearly that chemical equilibrium is reached already after ~10 days. The C–S–H behaviour in alkali-free conditions is assumed to be similar to that in ACW. The time evolution of the solution composition as observed in the present study is in contradiction with previous observations of L’Hôpital et al. (2015) who noticed a significant decrease in the aqueous Ca and Si concentrations during the first months of C–S–H ageing, indicating the formation of a more stable C–S–H phase with time. The final aqueous Ca and Si concentrations are however very similar in both studies. The slightly different procedures used during equilibration (horizontal shaking versus more vigorous end-over-end shaking used in the present synthesis procedure) might be a reason for the different kinetic behaviour as vigorous mixing of the suspensions could accelerate the ageing of amorphous phases.

All C–S–H phases used in the sorption studies in the following chapters were equilibrated for at least two weeks prior to the addition of the actinides or lanthanides under investigation.

Time evolution of the chemical composition of the solutions in contact with the other sorbents used in this study (TiO2, HCP) was not observed. In the case of HCP the material had been aged for at least 5 years prior to use in the sorption experiments. Thus, the C–S–H phases newly formed during the cement hydration were aged over a sufficiently long period of time to establish the most stable microstructural configuration.

Fig. 3: Evolution of the aqueous phase composition during the synthesis of C–S–H phases with three different C:S ratios in ACW at pH = 13.3. S:L = 0.02 kg L-1. a) Ca concentration. b) Si concentration

0 20 40 60 80 100 120 14010-5

10-4

10-3

10-2

C:S = 0.86 C:S = 1.15 C:S = 1.29Ca

conc

entra

tion

in so

lutio

n (M

)

Time (days)0 50 100 15010-5

10-4

10-3

10-2

C:S = 0.86 C:S = 1.15 C:S = 1.29

Si co

ncen

tratio

n in

solu

tion

(M)

Time (days)


4.2 Solid phase and aqueous phase composition

The composition of the C–S–H phases synthesized in this study was discussed in detail by Tits et al.(2006; 2014a) C–S–H phases with C:S ratios > 1.07 and synthesized in ACW contained significant amounts of portlandite whereas C–S–H phases with C:S ratios <≤ 1.07, were found to be portlandite-free. In alkali-free systems, portlandite was only detected in the C–S–H phases with a target C:S ratio of 1.82 although the aqueous Ca concentrations are close to the portlandite solubility in all C–S–H systems with C:S ratios > 1.29 (Fig. 3). Note, however, that the detection method for portlandite (XRD) is not at all sensitive and small quantities present at lower C:S ratios may have been overlooked. The portlandite contents and the actual C:S ratios corrected for portlandite are listed in Table 1. All C–S–H phases synthesized in ACW had significant amounts of Na and K bound to C–S–H . Alkali incorporation reached a maximum at the lowest C:S ratios and decreased with increasing C:S ratio.

Table 1: Bulk composition of selected C–S–H phases with different C:S ratios synthesized in ACW at pH = 13.3 and in MilliQ water

Target C:S ratio

mol/mol

Medium Ca(OH)2 content

1(Wt %)

Actual C:S ratio

(mol/mol)

2Na

mol kg-1

2K

mol kg-1

pH

0.64 0.75 1.07 1.29 1.50 1.82

ACW

0 0

<1 5.4

12.4 18.4

0.64±0.01 0.75±0.01 1.07±0.02 1.15±0.04 1.18±0.07 1.29±0.10

1.7±0.2 1.9±0.2 1.0±0.2 1.2±0.2 0.8±0.1 0.5±0.1

2.8±0.3 3.2±0.3 1.5±0.2 1.8±0.2 1.2±0.2 0.8±0.1

13.3 13.3 13.3 13.3 13.3 13.3

0.75 1.07 1.29 1.5 1.82

H2O

0 0 0 0

7.7

0.75±0.01 1.07±0.02 1.29±0.02 1.50±0.03 1.65±0.06

10.1 12.0 12.2 12.4 12.5

1Estimated uncertainties on the measurements is ±20%. 2Calculated per unit weight of CaO and SiO2 in the C–S–H phase

With increasing C:S ratio an increase of the Ca solubility and a decrease of the Si solubility was observed (Fig. 4) both in ACW at pH = 13.3 and in alkali-free conditions, in agreement with literature data (e.g. Chen et al., 2004; Lothenbach and Nonat, 2015 and references therein). In the absence of alkalis, the Ca concentration reaches a maximum value of ~2·10-2 M at a C:S ratio of ~1.65. This concentration corresponds to the Ca solubility w.r.t. portlandite in water. The Ca solubility of C–S–H phases synthesized in ACW at pH 13.3 is approximately a factor 10 lower than that of C–S–H phases with the same C:S ratio synthesized in alkali-free solutions. In this medium, the portlandite solubility of 1.6·10-3 M is reached at a target C:S ratio of 1.29 suggesting that a further increase in the target C:S ratio during synthesis does not produce C–S–H phases with higher effective C:S ratios but rather causes Ca(OH)2 precipitation. The Si solubility however, is quite similar in both media. The pH of the C–S–H suspensions in the absence of alkalis was found to correlate with the C:S ratio of the C–S–H phases (Table 1) and varied between 10.1 and 12.5. A detailed discussion of the composition of HCP used in the present report can be found in a paper by Lothenbach and Wieland (2006).


Fig. 4: Ca (a) and Si(b) solubility of C–S–H phases as a function of the actual C:S ratio in alkali-free solutions and in ACW at a constant pH of 13.3. Each datapoint represents an individual C–S–H phase. C:S ratios were calculated from mass balance calculations. The effective C:S ratio of the C–S–H phases with target C:S ratios ≥ 1.5 (in alkali-free solution) and ≥ 1.2 (in ACW) was corrected for the portlandite content based on the data in Table 1.

4.3 45Ca uptake by C–S–H phases

45Ca uptake experiments were carried out with C–S–H phases having a C:S ratio of 1.07 in alkali-free conditions and in ACW. In Fig. 5, the activity ratios (Asol/Atot) are plotted as a function of the equilibration time. The experimental data clearly show a two-step process: Very rapid uptake during the first 24 hours is followed by a much slower uptake process which is still not completed after 90 days. The fast uptake step is generally attributed to a reversible, linear surface adsorption process and is responsible for 50% to 80% of the 45Ca uptake in the present experiments. The second, slower step, however, represents 45Ca being incorporated in the C–S–H phase via recrystallization. This recrystallization process is still not completed after 90 to 120 days equilibration. The observed two-step uptake behaviour is consistent with earlier observations reported in the literature (Shrivastava et al., 1991; Baur et al., 2004; Mandaliev et al., 2010b). 45Ca uptake experiments carried out by Baur et al. (2004) showed that the rapid uptake step was less distinct but still present in experiments with pre-equilibrated C–S–H phases. This confirms the generally accepted concept that this rapid uptake step is caused by surface adsorption rather than recrystallization processes.

The rapid uptake surface adsorption step can be described by an equation similar to Eq. (3.4):

adsorbd

sol

A (t) VKA (t) m

= ⋅ , (4.1)

where Kd is the distribution coefficient resulting from the reversible, linear surface adsorption process (in contrast to Rd which describes the overall distribution ratio measured in the system at time t). The Kd value is independent of the reaction time whereas Aadsorb(t) and Asol(t) can vary with time due to the progressing 45Ca incorporation.

An homogeneous incorporation model was applied to account for the slow 45Ca incorporation process (Curti et al., 2010). During homogeneous incorporation, the entire recrystallized solid is

0.6 0.8 1.0 1.2 1.4 1.6 1.810-5

10-4

10-3

10-2

10-1

Alkali-free systems ACW

Solubility limit for Portlandite in ACW at pH 13.3

Ca eq

uilib

rium

conc

entra

tion

(M)

C:S ratio

Solubility limit for Portlandite at pH 12.5a

0.6 0.8 1.0 1.2 1.4 1.6 1.810-6

10-5

10-4

10-3

10-2

10-1

Solubility limit for SiO2(am) at pH 10.0

Alkali-free systems ACWSi

equi

libriu

m co

ncen

tratio

n (M

)

C:S ratio

b


at all times in equilibrium with the liquid phase. This model assumes that tracer diffusion takes place in the recrystallized solid from older recrystallized layers with higher 45Ca concentrations to more recent recrystallized layers with lower 45Ca concentrations. This results in a homogeneous distribution of the 45Ca tracer within the recrystallized C–S–H phase and, at all times, the ratio of radioactive and stable isotopes in the recrystallized solid equals the isotope ratio in solution (Curti et al., 2010):

45 45

recryst

recryst

Ca [ Ca]Ca [Ca]

= . (4.2)

Expressed in terms of activities, Eq. (4.2) can be written as:

incorp sol

recryst

A ACa [Ca] V

=⋅

. (4.3)

[Ca] and [45Ca] are the stable Ca and 45Ca concentrations in solution (M) and 45Carecryst and Carecryst are the moles of 45Ca and stable Ca in the recrystallized solid C–S–H phase. Aincorp and Asol are the 45Ca activities (Bq) in the recrystallized solid and the liquid phase, respectively, and V is the volume of the aqueous phase (L).

Fast tracer diffusion in the recrystallized C–S–H phases is a realistic assumption in view of the morphological models of C–S–H phases proposed recently in the literature: In these models C–S–H phases are described as an assembly of globules consisting of an agglomeration of small tobermorite-like or jennite-like particles with a diameter of a few nm’s (Thomas et al., 1998; Thomas and Jennings, 2006; Jennings, 2008; Chiang et al., 2012). In HCP, these globules are more or less densely condensed to form high density (HD) or low density (LD) C–S–H phases. Water-filled pores exist within the globules (interlayer space), in between the tobermorite-like particles, as well as in between the globules (gel pores). The gel-like properties and especially the high porosity of C–S–H phases enable optimal conditions for 45Ca tracer diffusion within the recrystallized C–S–H material.

Assuming a constant recrystallization rate, R (mol m-2 s-1), Carecryst can be expressed as:

recrystCa R t SA m= ⋅ ⋅ ⋅ , (mol) (4.4)

R is the recrystallization rate (mol m-2 s-1), SA is the specific surface area (m2 kg-1), and m is the amount of solid (kg).

The following equations were derived to account for the two-step uptake kinetics (surface sorption + homogeneous incorporation):

tot C S H solA A (t) A (t)− −= + , (Bq) (4.5)

C S H adsorb incorpA (t) A (t) A (t)− − = + , (Bq) (4.6)

with Atot, AC–S–H (t), Aadsorb(t) and Aincorp(t) as the total 45Ca activity in the system, the total 45Ca activity (adsorbed + incorporated) in the C–S–H phase, the 45Ca activity adsorbed onto the C–S–H surface and the 45Ca activity incorporated in the C–S–H structure at time t, respectively.

Aincorp(t) can be derived by combining Eq. (4.3) and (4.4), and Aadsorb(t) can be derived from Eq. (4.1).


Combining Eq. (4.1) to (4.6) results in the following expression for the activity ratio in solution for a homogeneous incorporation model combined with reversible surface adsorption:

sol

tot d

A (t) [Ca]A (t) K (m / V) [Ca] R (m / V) SA t [Ca]

=⋅ ⋅ + ⋅ ⋅ ⋅ +

. (4.7)

In this equation R and Kd are the fitting parameters. The specific surface area was fixed to 1.48·105 m2 kg-1 (Tits et al., 2006). The fitting results show that the homogeneous incorporation model combined with rapid surface adsorption reasonably describes the time evolution of 45Ca uptake by this C–S–H phase (Fig. 5). The fitted R values were estimated at 2·10-12 mol m-2 s-1 and 4·10-12 mol m-2 s-1 for the C–S–H phase in alkali-free solution and in ACW, respectively. The Kd values were determined to be 500 L kg-1 in alkali-free system and 1800 L kg-1 in ACW.

Fig. 5: 45Ca activity ratios for two C–S–H recrystallization experiments (S:L ratio = 2·10-3 kg L-1) in a) alkali-free solution and b) in ACW. Red lines in (a) and (b) are fits with a model combining homogeneous incorporation with surface sorption. c) fraction of C–S–H recrystallized

Knowing the recrystallization rate R, the amount of C–S–H recrystallized can be calculated as a function of time using Eq. (4.4). Multiplying Carecryst with the molecular weight (mW) of the C–S–H phase and dividing the term by the mass of C–S–H (mC–S–H ) results in the mol fraction of C–S–H phase recrystallized:

0 30 60 90 120 1500.0

0.2

0.4

0.6

0.8

1.0

Experimental data Homogeneous incorporation + sorption

Aso

l/Ato

t

Time (days)

a Alkali-free systems

0 30 60 90 120 150

0.00

0.05

0.10

0.15

0.20

0.25

0.30

Experimental data Homogeneous incorporation + surface sorption

Aso

l/Ato

t

Time (days)

b ACW

0 30 60 90 120 1500.0

0.2

0.4

0.6

0.8

1.0 ACW Alkali-free

Frac

tion

recr

ysta

llize

d C

-S-H

Time (days)

c


CSHrecryst recryst w C S HF Ca m / m − −= ⋅ . (4.8)

mWC–S–H is defined in Eq. (2.1) (Section 2.1.2).

The model calculations shown in Fig. 5c indicate that in the conditions of the recrystallization experiments reported here, 40% to 67% of the C–S–H phases were recrystallized within 150 days equilibration.

It is to be noted that recrystallization rates may depend on the specific experimental conditions, such as crystallinity of the phases (ageing), S:L ratio, etc. The recrystallization rates obtained from the recrystallization experiments described in this section cannot be applied directly to the C–S–H systems in the sorption studies presented in the following sections. They only illustrate that actinide incorporation by recrystallization of the C–S–H phases with time is a perfectly feasible process.

It is presently uncertain whether or not the current recrystallization rates also hold for C–S–H phases which form in HCP during the hydration process and subsequently age with time. The experiments with synthesized C–S–H phases clearly show that, in principle, recrystallization of the C–S–H phases in HCP should occur while the rate might be significantly lower.

5 Sorption of trivalent actinides and lanthanides

5.1 Speciation

The most important trivalent actinide relevant for low and intermediate level radioactive waste is Am(III). Eu(III) is often used as a chemical analogue for Am(III) because its stable isotopes are easier to handle in the laboratory and allow sorption experiments to be conducted over a very large concentration range. Comparison of the complexation constants of Eu(III) and Am(III) (Table 2), however, shows that the stabilities of Eu(III) and Am(III) complexes with the same ligands can differ significantly. The experiments described in the present report are mainly performed with Eu(III) because of the above-mentioned practical reasons. Only in a few cases, parallel sorption experiments with trivalent actinides (Cm(III) and Am(III)) were performed. In all these cases, Rd values for trivalent actinide sorption were found to be fairly similar (Am(III)) to, or even higher (Cm(III)) than the Rd values determined for Eu(III).

In view of the composition of C–S–H pore waters, the aqueous trivalent actinide/lanthanide speciation will mainly consist of hydroxy complexes and eventually ternary actinide/lanthanide-Ca-hydroxy complexes. In addition, actinide/lanthanide–silicate and ternary actinide/lanthanide hydroxyl silicate complexes may form. However, although such species probably exist, the lack of thermodynamic data for these species in the literature does not allow making statements about their presence. The speciation calculations presented below were carried out with the code “Medusa” (Puigdomenech, 1983) using the thermodynamic complexation constants for Am(III) and Cm(III) listed in the NEA thermodynamic database (Guillaumont et al., 2003) and completed with complexation constants measured for Cm(III) as given in Neck et al. (2009) and Rabung et al. (2008). The speciation of the trivalent lanthanide (Eu(III)) was calculated using the thermodynamic complexation constants for Nd(III) reported by Herm et al. (2015) and Neck et al. (2009) and is identical to the trivalent actinide speciation under alkaline conditions. The thermodynamic data for Eu(III) reported by Hummel et al. (2002) were not used for data interpretation because they significantly overestimate the predominance of the Eu(OH)4

- species


at high pH in comparison with speciation data for Nd(III) and the trivalent actinides published in the literature (Table 2). To date, the formation of Am(III)-silicate, Cm(III)-silicate and Eu(III)-silicate complexes has only been investigated at low pH (pH < 5). The complexes observed in the acid pH region (AmSiO(OH)3

2+, CmSiO(OH)32+ and EuSiO(OH)3

2+), however, are not relevant for the actinide/lanthanide speciation under alkaline conditions. Data on Am(III)-silicate or Cm(III)-silicate complexation in neutral and alkaline media are completely lacking in the literature. Therefore, we cannot account for these complexes in the thermodynamic calculations although they might play an important role in some C–S–H pore waters (pH = 10, high Si concentrations). A summary of the known relevant thermodynamic complexation constants of Am(III)/Cm(III), Nd(III) and Eu(III) is given in Table 2. Constants for all other reactions needed in the calculations were taken from the core dataset published by Hummel et al. (2002). All calculations were performed fixing I to 0.3 M. This value reflects the ionic strength in ACW. Note that this implies that the speciation calculated at pH > 13.3 is only approximately correct.

Speciation calculations were performed for two different chemical systems: 1) a system containing no Ca and a system containing 0.02 M Ca. In the latter system the concentration of the aqueous Ca species decreases suddenly above pH = 12.4 due to portlandite precipitation (not shown). The calculations displayed in Fig. 6 reveal clear differences in the speciation of the trivalent actinides and Nd(III) in contrast to Eu(III): In the case of Am(III) and Nd(III), the speciation in the alkaline pH range is dominated by the Am(OH)2

+ and the Am(OH)3 hydroxy complexes. Neither the Am(OH)4

- complex, nor the ternary Cap(Am(OH)n)2p+3-n complexes appear in cementitious pore waters. When the thermodynamic data for Eu(III) were used, the Eu(OH)3 and Eu(OH)4

- complexes clearly dominate the speciation, which is rather unrealistic in view of the data for Nd(III) and thos of the trivalent actinides reported in the literature.

Table 2: Relevant thermodynamic complexation constants for trivalent actinides and lanthanides used in the speciation calculations presented in Fig. 6 (Hummel et al., 2002; Guillaumont et al., 2003; Neck et al., 2009; Thoenen et al., 2014)

Hydroxide complexes: log*β0n: 3 3 n

2 nAn n H O An(OH) n H+ − ++ ⋅ ⇔ + ⋅ Am(III)/Cm(III) Eu(III)

AnOH2+ An(OH)2

+ An(OH)3 An(OH)4

-

-7.2 -15.1 -26.2 -40.7

EuOH2+ Eu(OH)2

+ Eu(OH)3 Eu(OH)4

-

-7.64 -15.1 -23.7 -36.2

NdOH2+ Nd(OH)2

+ Nd(OH)3 Nd(OH)4

-

-7.4 -15.7 -26.2 -40.7

Ternary complexes: log*β0p,1,n: 2 3 2p 3 n

2 p npCa An n H O Ca [An(OH) ] n H+ + + − ++ + ⋅ ⇔ + ⋅ CaAn(OH)3

2+

Ca2An(OH)43+

Ca3An(OH)63+

-26.3 -37.2 -60.7

Solubility: log*K0s: 3

3 2An(OH) (cr) 3H An 3H O+ ++ ⇔ + An(OH)3(am) 16.9 Eu(OH)3(am) 17.6 Nd(OH)3(am) 17.2


Fig. 6: Trivalent actinide and lanthanide speciation as a function of pH in the absence of carbonates. The total Am(III) and Eu(III) concentrations were fixed at 10-5 M. a and b) Am(III) speciation (and Nd(III) speciation) in absence of Ca and in the presence of 2·10-2 M Ca, c) Eu(III) speciation. Thermodynamic equilibrium constants taken from Table 2.

5.2 Preliminary solubility experiments

The stability of Eu(III) solutions was only tested in ACW in the concentration range between 10-11 M and 10-7 M. The stability of Am(III) solutions was not tested. Based on the speciation calculations shown in Fig. 6a and 6b, the solubility of Eu(III) is expected to decrease with increasing pH. Unfortunately solubility tests in alkali-free solutions were not carried out. The results of the solubility tests are presented in Fig. 7. Activity losses between 40% and 80% were observed over the entire concentration range before centrifugation, due to sorption on the container walls (Figs. 7a and 7c). The fractions of the Eu(III) inventories measured in solution before centrifugation were lower after 30 days than after 1 day, suggesting that sorption on the container walls was a slow process. Centrifugation (1 h at 95,000 g (max)) caused additional losses at radionuclide concentrations above ~10-9 M. This loss of activity is explained by the presence of Eu(III) colloids which were generated in ACW and settled during centrifugation. The data displayed in Figs. 7b and 7d show the equilibrium Eu(III) concentration as function of the input concentration. The solid lines with slope 1 represent the case where the input concentration is completely recovered in the equilibrium solution. Both datasets determined before centrifugation lie below this line and approximately parallel to it, indicating sorption of

10 11 12 13 14-12

-11

-10

-9

-8

-7

-6

Am(OH)3(s) solubility

Am3+

AmOH2+Am(OH)-

4

Am(OH)+2

log

conc

entra

tion

(M)

pH

Am(OH)3

[Ca] = 0 Ma

10 11 12 13 14-12

-11

-10

-9

-8

-7

-6

Am3+

Am(OH)3(s) solubility

Ca(Am(OH)3)+2

AmOH2+

Am(OH)-4

Am(OH)+2

log

conc

entra

tion

(M)

pH

Am(OH)3

[Ca] = 0.02 Mb

10 11 12 13 14-12-11-10-9-8-7-6-5-4-3

Eu(OH)3(s) solubility

Eu3+

EuOH2+

Eu(OH)-4

Eu(OH)+2

log

conc

entra

tion

(M)

pH

Eu(OH)3

c [Ca] = 0 M


the radionuclides on the container walls. The presence of Eu(III) radiocolloids was observed at aqueous concentrations above approximately 10-9 M which is somewhat higher than the calculated solubility limit for Nd(OH)3(am) represented by the dashed horizontal line in Figs. 7b and 7d. The solubility threshold observed for the Eu(III) solutions in ACW is much lower than the thermodynamic solubility limits calculated for Eu(III) (Figs. 6c), confirming our conclusion that the Eu(III) thermodynamic data listed in the thermodynamic database of Hummel et al. (2002) overestimate the dominance of the aqueous Eu(OH)3 and Eu(OH)4

- complexes and the Eu(III) speciation is very similar to that of Am(III) and Nd(III).

Fig. 7: Eu(III) concentration measured in solution before and after centrifugation as a function of the input Eu(III) concentration. Measurements were conducted after 1 day (a, b) and 30 days equilibration (c, d). The solid lines in b and d represent 100% Eu(III) recovery in solution. The dashed horizontal lines represent the calculated Eu(OH)3(am) solubility limit in ACW based on the thermodynamic data for Nd(III). Dotted lines are added to guide the eye.

5.3 Eu(III) sorption and co-precipitation kinetics

5.3.1 C–S–H phases Eu(III) sorption kinetics were studied with C–S–H phases having C:S ratios varying between 0.75 and 1.5 in ACW (Fig. 8a). Sorption kinetic experiments in alkali-free conditions were not

10-11 10-10 10-9 10-8 10-70

20

40

60

80

100 before centrifugation after centrifugation

% o

f add

ed E

u(III

) in

solu

tion

Concentration of added Eu(III) (M)

1 day equilibration

a

10-11 10-10 10-9 10-8 10-7 10-6

10-11

10-10

10-9

10-8

10-7

10-6

10-5

Experimental Eu(III) stability limit

1 day equilibration before centrifugation after centrifugation

Eu(OH)3(s)

Eu(II

I) co

ncen

tratio

n in

solu

tion(

M)


b

10-11 10-10 10-9 10-8 10-7 10-6

0

20

40

60

80

100

30 days equilibration

before centrifugation after centrifugation

% o

f add

ed E

u(III

) in

solu

tion

Concentration of added Eu (M)

c

10-11 10-10 10-9 10-8 10-7 10-610-12

10-11

10-10

10-9

10-8

10-7

10-6

10-5


Eu(OH)3(s)

before centrifugation after centrifugation

Eu(II

I) co

ncen

tratio

n in

solu

tion(

M)


d

Experimental Eu(III) stability limit


performed. The sorption experiments were carried out at C–S–H concentrations (S:L ratios, based on the dry weight) of approximately 5·10-4 kg L-1 and the total Eu(III) concentration varying between ~2·10-9 M and ~10-8 M. The maximal Rd value measureable in these experiments, Rd,max, was estimated to be 1.7·107 L kg-1, which is significantly above the measured Rd values shown in Fig. 8. The results reveal a fast sorption kinetic behaviour. Steady state was reached within 1 day. Rd values were found to remain constant over a time period of up to 90 days. Very high Eu(III) distribution ratios (Rd = (7±4)·105 L kg-1) were determined independent of the C–S–H composition (C:S ratio). The drying step during the preparation of the C–S–H phases, which was applied in an earlier series of C–S–H synthesis, did not have any significant influence on the Rd values (red symbols in Fig. 8).

A second type of kinetic studies was devoted to the co-precipitation of Eu(III) with C–S–H phases in ACW. In this type of experiments, 152Eu tracer was added to a suspension of CaO and SiO2 in ACW during C–S–H synthesis and the uptake of this tracer by the precipitating C–S–H phase was determined as a function of time. In Fig. 8b, the data obtained from the co-precipitation experiments are presented in terms of Rd values as a function of the reaction time. The extent of Eu(III) uptake in these experiments is very similar to the uptake observed in the Eu(III) sorption experiments with C–S–H phases in ACW: i.e. very fast uptake and high distribution ratios. This indicates rapid and complete removal of the Eu(III) from solution. From a statistical point of view, the Rd values obtained from the co-precipitation experiments seem to be slightly lower than those obtained from the sorption experiments (Fig. 8b). Nevertheless, there is no reason to assume that co-precipitation could lead to lower Rd values. For this reason and also because of the large uncertainties on the measurements, the Rd values for Eu(III) sorption and Eu(III) co-precipitation with C–S–H phases are considered to be in good agreement. The absence of a long term kinetic effect on the Rd value observed in the course of the sorption and co-precipitation experiments with C–S–H phases suggests that recrystallization of these cementitious materials does not have a significant influence on the Eu(III) uptake. These observations seem to contradict the abundant spectroscopic evidence for incorporation of trivalent lanthanides in C–S–H phases. (Pointeau et al., 2001; Tits et al., 2003; Schlegel et al., 2004; Stumpf et al., 2004; Macé et al., 2007; Mandaliev et al., 2010a; 2010b).

Fig. 8: a) Eu(III) sorption kinetics onto C–S–H phases with different C:S ratios in ACW at pH = 13.3. Experimental conditions: S:L = 5·10-4 kg L-1, 1.98·10-9 M < [Eu]tot < 9.9·10-9 M. Red symbols: experiments carried out with dried C–S–H phases. Black symbols: experiments carried out with fresh C–S–H suspensions. b) Eu(III) – C–S–H co-precipitation kinetics in ACW at pH = 13.3. Comparison with the data from the sorption kinetic experiments (datapoints in grey). The dashed lines represent the Rd,max value.

0 20 40 60 80 100101

102

103

104

105

106

107

C:S = 0.75 C:S = 0.95; exp. 1 C:S = 0.95; exp. 2 C:S = 1.07 C:S = 1.2; exp. 1 C:S = 1.2; exp. 2 C:S = 1.4 C:S = 1.5; exp. 1 C:S = 1.5; exp. 2 Mean

R d (E

u(III

)) (L

kg-1

)

Time (days)

ACWa

0 20 40 60 80 100101

102

103

104

105

106

107

C:S = 0.75 C:S = 1.25 Mean Rd

R d (E

u(III

)) (L

kg-1

)

Time (days)

ACWb


5.3.2 Hardened cement paste 152Eu(III) sorption kinetic experiments with HCP in ACWHCP-I and ACWHCP-II confirm the general picture observed for Eu(III) sorption onto C–S–H phases (Fig. 9). Fast sorption kinetics leads to a steady state within one day and very strong sorption characterized by very high Rd values of (2±1)·105 L kg-1. Sorption kinetic tests with Am(III) onto HCP show a sorption behaviour onto HCP similar to Eu(III) notwithstanding differences in speciation (Fig. 6). i.e. fast sorption kinetics and very strong sorption illustrated by Rd values ranging between 5·104 L kg-1 and 5·105 L kg-1.

Fig. 9: Eu(III) sorption kinetics (a) and Am(III) sorption kinetics (b) onto HCP in ACWHCP-I (pH = 13.3) and ACWHCP-II (pH = 12.5). Experimental conditions: S:L = 10-4 kg L-1, [Eu]tot = 2·10-9 M, [Am]tot = 3·10-10 M.

The apparent positive effect of pH on the Am(III) Rd values is currently not understood and might be due to an experimental artefact as the aqueous Am(III) speciation calculations cannot provide any justification for the trend. On the contrary, higher pH values would rather stabilize aqueous Am-hydroxy complexes or aqueous Am-silicate complexes (if they existed) and thus weaken Am(III) sorption onto HCP.

5.4 Desorption tests

Eu(III) desorption tests were carried out as follows: After sorbing Eu(III) for a specific period of time, the solid and liquid phases were separated by centrifugation and the liquid phase was replaced by a fresh solution having a composition identical to the original liquid phase but without Eu(III). The 152Eu activity released from the C–S–H phase into the fresh solution was determined as a function of time. The Rd,max value for these experiments is similar to the value deduced for the sorption kinetic tests (Section 5.3.1.). Fig. 10 shows the results of two desorption experiments with C–S–H phases having C:S ratios of 0.75 and 1.25 in ACW. The experimental data show that distribution ratios (Rd values) obtained from the desorption experiments after one day sorption and after 60 days sorption agree well with those determined in the original sorption tests indicating that Eu(III) sorption on the C–S–H phases, even after 60 days sorption, is still reversible.

0 20 40 60 80 100104

105

106

107

ACW-I; pH = 13.3 ACW-I; pH = 13.3 ACW-II; pH = 12.5 ACW-II; pH = 12.5

R d (E

u(III

)) (L

kg-1

)

Time (days)

a Eu(III)

0 50 100 150 200 250 300 350 400104

105

106

107

ACW-I; pH = 13.3 ACW-I; pH = 13.3 ACW-II; pH = 12.5 ACW-II; pH = 12.5

R d (A

m(II

I)) (L

kg-1

)

Time (days)

b Am(III)


Fig. 10: Eu(III) desorption tests onto C–S–H phases in ACW at pH = 13.3 after one day and 60 days sorption. a) C:S = 0.75; b) C:S = 1.25. Experimental conditions: S:L = 2.5·10-3 kg L-1, [Eu]tot = 3.0·10-8 M. The dashed lines represent the Rd,max value.

5.5 Effect of the S:L ratio

Rd values for Eu(III) sorption onto C–S–H phases, measured in suspensions with varying S:L ratios, are shown in Fig. 11a. The Rd,max values for these experiments decrease with increasing S:L ratio (dashed line in Fig. 11a). The sorption data displayed in Fig. 11 show a slight decrease in the Rd value from ~106 L kg-1 at a S:L ratio of 10-5 kg L-1 to ~2·105 L kg-1 at a S:L ratio of 2·10-2 kg L-1. At first glance, this effect of the S:L ratio seems to be experimentally significant. However, taking into account the large scatter on the experimental data in this experiment and even more in the kinetic experiments shown in Fig. 7, it may be challenged whether or not the trend is real. At present, it is difficult to reach a final conclusion.

A similar sorption behaviour was observed for Eu(III) onto HCP (Fig. 11b). In this case sorption values decreased from ~8·105 L kg-1 at a S:L ratio of 7·10-5 kg L-1 to ~4·104 L kg-1 at a S:L ratio of 2·10-2 L kg-1. The experimental data measured at a S:L ratio of 10-5 kg L-1 are assumed to be artefacts. Addition of 1.0 M NaNO3 to the suspensions prior to centrifugation, however, improved phase separation via enhanced colloid flocculation which gave rise to higher Rd values at the highest S:L ratios. As a consequence, the S:L ratio dependence of the Eu(III) sorption was much weaker. A similar influence of the S:L ratio was reported in the literature for Eu(III) sorption onto calcite (Tits et al., 2005).

The presence of sorption sites with different affinity for Eu(III) cannot explain the observed dependence of the Rd value on the S:L ratio. Lower S:L ratios would provoke saturation of higher affinity sorption sites so that part of the sorbing Eu(III) is forced to occupy sorption sites with a lower affinity, thus reducing the overall Rd value. By contrast, the S:L dependent sorption behaviour of Eu(III) sorption reveals a trend in the opposite direction. The higher the concentration of the high affinity sites (increasing S:L ratio), the lower the Rd value. The experiments carried out by adding high concentrations of NaNO3 before phase separation suggest that the observed effect is due to incomplete removal of colloidal material loaded with 152Eu during centrifugation. The observed effects of the S:L ratio on the Eu(III) Rd values are currently not understood.

0 10 20 30 40 50 60 70 80 90 100102

103

104

105

106

107

Rd sorption Rd desorption after 1 day sorption Rd desorption after 60 days sorption

C:S = 0.75

R d (E

u(III

)) (L

kg-1

)

Reaction time (days)

a

0 10 20 30 40 50 60 70 80 90 100102

103

104

105

106

107C:S = 1.25

Rd sorption Rd desorption after 1 day sorption Rd sorption Rd desorption after 60 days sorption

Rd (

Eu(I

II))

(L k

g-1)


b


Fig. 11: Eu(III) sorption onto C-S-H phases and HCP as function of the S:L ratio.

a) C–S–H phase (C:S = 1.07) in ACW at pH = 13.3. Experimental conditions: equilibration time = 2 weeks, [Eu]tot = 8.5·10-10 M. The dashed line corresponds to the Rd,max values as function of the S:L ratio. b) Fresh HCP in ACW-I at pH 13.3: equilibration time = 2 weeks, [Eu]tot = 7.2·10-10 M.

5.6 Sorption isotherms

An Eu(III) sorption isotherm was determined in ACW for a C–S–H phase with a C:S ratio of 1.0 at a S:L ratio of 5·10-4 kg L-1 and an equilibration time of 2 weeks. The total Eu(III) concentration in the experiments was varied between 10-10 M and ~5·10-8 M. The data are shown in Fig. 12. The Rd,max value (dashed line in Fig. 12) increases with increasing total Eu(III) concentration in the systems because the lower part of the isotherm was obtained by adding increasing concentrations of 152Eu, whereas in the higher part, increasing concentrations of stable Eu(III) labelled with a fixed 152Eu activity were added. The isotherm was found to be linear over the entire investigated concentration range (slope = 1.05±0.06). The highest Eu(III) loading covered by the sorption isotherm is 2·10-4 mol kg-1 or 6·10-4 eq kg-1 (assuming that one Eu(III) neutralizes three sorption sites on the C–S–H phase). This Eu(III) loading corresponds to ~0.01% of the sorption capacity of the C–S–H phases (~6 eq kg-1) or also ~0.01 % of the amount of Ca2+ available for substitution in the CaO layer. This low loading implies that only the sorption sites with the highest affinity for Eu(III) are occupied which explains the linear sorption behaviour of Eu(III). The mean Rd value calculated from the measured data on the entire concentration range was (4±2)·105 L kg-1 which, within the given uncertainty of the sorption data, agrees with the sorption kinetic data (Fig. 6) and the sorption data obtained as function of the S:L ratio (Fig. 11).

10-5 10-4 10-3 10-2104

105

106

107

108 Exp. 1 Exp. 2 Exp. 3

R d(E

u(III

)) (L

kg-1

)

S:L ratio (kg L-1)

a

10-5 10-4 10-3 10-2 10-1104

105

106

107

108

( )

ACW-I; pH = 13.3; exp 1 ACW-I; pH = 13.3; exp. 2 ACW-I; pH = 13.3; NaNO3

R d (E

u(III

)) (L

kg-1

)

S:L ratio (kg L-1)

b

( )


Fig. 12 a,b) Eu(III) sorption isotherm in ACW onto a C–S–H phase with C:S ratio = 1.0. Experimental conditions: S:L = 5·10-4 kg L-1, equilibration time = 2 weeks. Cm(III) sorption experiments: S:L = 10-3 kg L-1. c,d) Eu(III) and Am(III) sorption isotherms onto fresh HCP. Experimental conditions: S:L = 10-4 kg L-1; equilibration time = 2 weeks. The dashed line in a), represents Rd,max values for Eu(III).

Eu(III) and Am(III) sorption isotherms onto HCP (Figs. 12c and 12d) confirm the behaviour previously observed on C–S–H: linear sorption over a large concentration range up to an equilibrium concentration of ~2·10-7 M. In the case of Eu(III) a further increase of the added Eu(III) concentration results in precipitation of Eu(OH)3(s) colloids. The observed solubility threshold is higher than the calculated Am(OH)3(s) / Nd(OH)3(s) solubility limit (~8·10-6 M, Fig. 6a, b) and also higher than the measured solubility limits for Eu(III) solutions in ACW (~10-8 M, Figs. 7a – 7d). This observation is currently not understood. The synthetic ACW used in the preliminary solubility tests in section 5.2 do not contain Si. It can therefore be speculated that the presence of Si in the HCP solution may have increased the Eu(III) solubility due to the formation of ternary Eu(III)-hydroxy-silicate complexes. The series of measurements conducted with Am(III) in the low concentration range confirm the identical sorption behaviour of Eu(III) and Am(III).

10-13 10-12 10-11 10-10 10-9104

105

106

107

108

Eu(III) Cm(III)

R d (L

kg-1

)

Equilibrium concentration (M)

a

ACW

10-13 10-12 10-11 10-10 10-910-8

10-7

10-6

10-5

10-4

10-3

10-2

Eu(III) Cm(III)

Eu(II

I) (so

rbed

) (m

ol k

g-1)


b

10-12 10-11 10-10 10-9 10-8 10-7 10-6103

104

105

106

107

108

HCP in ACWHCP-I at pH = 13.3

Eu(III) exp. 1 Eu(III) exp. 2 Am(III)R d

(An(

III)/L

n(III

)) (L

kg-1

)


c

10-12 10-11 10-10 10-9 10-8 10-7 10-610-7

10-6

10-5

10-4

10-3

10-2

10-1

100

Eu(III) exp. 1 Eu(III) exp. 2 Am(III)

An(II

I)/Ln

(III)

sorb

ed (M

)


d HCP in ACWHCP-I at pH = 13.3


5.7 Discussion

5.7.1 Sorption data The sorption behaviour of Eu(III) and Am(III) on C–S–H and HCP can be summarized as follows:

1. Similar sorption behaviour for Am(III) and Eu(III); 2. Strong sorption illustrated by Rd values in the range 105 L kg-1 < Rd < 106 L kg-1; 3. Fast and reversible uptake process with steady state reached within 1 day; 4. Linear sorption; 5. Indications for decreasing Rd values with increasing S:L ratios caused by experimental

constraints; 6. No effect of the C–S–H composition (C:S ratio) on the extent of sorption; 7. A slight effect of degradation of HCP (effect of pH?); 8. Rd values obtained from sorption experiments and from co-precipitation experiments agree

suggesting similar sorption mechanisms.

Thermodynamic calculations (Fig. 6) show that the Am(III) and Eu(III) speciation varies only little with pH in the range 10 < pH < 13.3, when the thermodynamic constants of Nd(III) are used for the Eu(III) speciation calculations. In both cases, the dihydroxy species dominates the speciation in the pH range between 10 and 11 and the trihydroxy species dominates the speciation above pH 11. The tetrahydroxy species plays a role only above pH = 13.5 and is not relevant in our experiments. The theory of the electrostatic interligand repulsion (section 2.3.) predicts that the limiting number (nlimit) of OH groups that fit in the first coordination sphere of trivalent actinides and lanthanides, is 4 and that the formation constants for the addition of a 3rd or a 4th OH group in the first coordination sphere of the metal cation, is similar (Fanghänel and Neck, 2002, Fig, 1b). Thus, this theory predicts that the sorption of dihydroxy- and trihydroxy Am(III) and Eu(III)species will be similar confirming the observations made on Figure 9 that sorption of Eu(III) and Am(III) on HCP at pH = 12.5 and 13.3 is similar. In the pH range 10 to 12.5 no experimental sorption data were collected.

Measured sorption values for trivalent actinides reported in the literature are generally lower than the Rd values for Eu(III) sorption presented in this work: Pointeau et al. (2004a) and Cowper et al. (2006) published Am(III) sorption data on C–S–H phases and HCP. Pointeau et al. (2004a) studied the uptake of Am(III) by C–S–H phases and HCP at different degradation stages in the pH range 9 < pH < 12. They observed high Rd values ranging from ~2·104 L kg-1 to ~2·105 L kg-1. Note that the authors did not specify the experimental “window” in particular the authors did not estimate the Rd,max value on the basis of their experimental set-up. Note further that the upper limit of the range of Rd values agrees with the data presented in this report.

In the work of Pointeau et al. (2004a) the Rd values tended to be independent of pH in the pH range under investigation which agrees with the above outlined expectations of pH effects on Eu(III) and Am(III) sorption. In contrast to the present sorption data, Pointeau et al. (2004a) observed that the Am(III) sorption process onto degraded HCP was irreversible as evidenced by higher Rd values measured in desorption tests. Cowper et al. (2006) reported much weaker sorption onto fresh HCP with Rd values ranging between 3·103 L kg-1 < Rd < 7·103 L kg-1 (pH = 12.9). These authors further report higher sorption values (~105 L kg-1) onto degraded HCP at lower pH (11.6). Also these authors did not discuss the experimental “window” relevant for their study; in particular the authors did not address the maximum Rd value measurable in the cementitious systems used.


5.7.2 Uptake mechanisms In Section 2.1.2, silanediol sites in the C–S–H interlayers and silanol sites at the edges of C–S–H phases were postulated as potential sites for actinide/lanthanide sorption. In addition, lanthanides and trivalent actinides can easily substitute for calcium ions in the Ca-O layers due to their similar ionic radii (rEu = 1.07Å, rNd = 1.11Å, rAm = 1.10Å, rCm = 1.09Å) as compared to the ionic radius of Ca2+ (rCa = 1.12Å) assuming coordination number CN = 8 (Shannon, 1976). Recently, several research groups have published results of spectroscopic studies indicating incorporation of trivalent actinides and lanthanides in C–S–H phases (Pointeau et al., 2001; Tits et al., 2003; Schlegel et al., 2004; 2009; Mandaliev et al., 2010a; 2010b; 2011). Several publications provide strong evidence for a two-step process (Tits et al., 2003; Mandaliev et al., 2010a; 2010b; 2011): in a first step the trivalent actinide or lanthanide rapidly forms an inner surface complex or a surface precipitate on the C–S–H phases. This step is followed by a slower second process during which the trivalent actinide or lanthanide becomes incorporated in the C–S–H structure, substituting for Ca2+ in the C–S–H interlayer and in the Ca-octahedral layer.

The existence of such a two-step process was clearly demonstrated in the fluorescence emission spectra of Cm(III) sorbed on a C–S–H phase as a function of time (Fig.13; Tits et al., 2003). After 1 hour contact time, Cm(OH)3 formed a surface precipitate on the C–S–H phases. The Cm(OH)3 surface precipitate did not emit fluorescence light due to concentration quenching whereby the light from fluorescing Cm ions was absorbed by neighbouring non-excited Cm ions in the precipitate. With time, an increasing fraction of the Cm(III) became incorporated in the C–S–H structure which resulted in an increasing intensity of the fluorescence signal along with a red-shift of the Cm fluorescence band. The lifetime of the Cm fluorescence provides information on the composition of the first coordination sphere of the fluorescing Cm species. Kimura and Choppin (1994) showed that a linear correlation exists between the Cm fluorescence decay rate and the number of H2O molecules in the first coordination sphere of Cm. Using this correlation, Tits et al. (2003) determined two Cm sorbed species on C–S–H phases, a first species with 1.4 H2O molecules left in its first coordination sphere and a second species with no H2O molecules in its first coordination sphere. In the original publication, the authors proposed that the first sorbed Cm species substitutes for Ca2+ in the interlayer coordinating with seven oxygens, where five belong to silica tetrahedra and two belong to adjacent H2O molecules in the interlayer. However, such a coordination environment requires that the sorbed Cm loses its OH-groups upon sorption which would result in pH dependent Rd values. Such pH dependence was not observed in the batch sorption experiments discussed above. In a more recent interpretation of the Cm fluorescence decay, it was postulated that the quench rate of Cm is not determined by the number of H2O molecules but by the number of quenching OH- oscillators in the first coordination sphere (Supkowski and Horrocks, 2002). A H2O molecule contains two OH oscillators thus the quench rate of an OH-group is approximately half that of a H2O molecule. Instead of having 1.4 H2O molecules in its first coordination sphere, the lifetime of the first Cm sorbed species can then also be explained by the presence of 2.8 (2 x 1.4) OH groups, meaning that Cm(OH)3 species occupy the interlayer sorption sites instead of free Cm3+ cations. This new interpretation is consistent with the speciation calculations (Fig. 6) showing that neutral tri-hydroxy actinide and lanthanide species (An(OH)3 and Ln(OH)3) are dominant in the largest part of the pH range under investigation. Note that the observed intercalation of Cm(OH)3 in the C–S–H interlayers, apparently does not go with a release of Ca from interlayer. Indeed, an increase in aqueous Ca concentrations from ∼10-5 M to 1.6·10-3 M with increasing C:S ratio of the C–S–H composition, did not exert any effect on the Eu(III) Rd values. The higher aqueous Ca concentration would have reduced Eu(OH)3 intercalation significantly when Eu(OH)3 substitution for Ca in the interlayer would have been the driving process.


The second Cm sorbed species, having lost all its water molecules (or OH oscillators) in its first coordination sphere, is assumed to occupy structural Ca sites in the Ca-O layers of the C–S–H phase. Such a substitution process requires that the OH groups are removed from the Cm(III) coordination shell, a process that becomes increasingly difficult with increasing pH. Such a pH effect was not observed in the batch sorption experiments, but it is believed that any change in Rd value due to such a pH effect would be within the scatter of the experimental data. Furthermore, substitution of An(III) or Ln(III) for Ca in the Ca-O layer is in contradiction with the observation (Fig. 7) that the C:S ratio has no influence on the Rd values. Indeed, an increase in the C:S ratio from 0.67 to 1.65 in ACW results in an increase in the Ca2+ concentration from ∼10-5 M to 1.6·10-3 M. Such an increase would inevitably cause a reduction of the Rd values in the case of a substitution controlled An(III)/Ln(III) uptake process. This is obviously not the case. Hence, Ln(III)/An(III) incorporation into C–S–H phases seems to take place without significant release of Ca.

Fig. 13: Fluorescence emission spectra of 10-7 M Cm(III) sorbed onto a C–S–H phase (C:S = 1.07) in ACW at pH 13.3. a) Fluorescence emission as function of the reaction time. b) Fluorescence emission spectra after 58 days contact time recorded at increasing delay times. The time dependence of the emission decay shows a bi-exponential decay behaviour. A detailed discussion of the Cm(III) TRLFS experiments is given in the paper of Tits et al. (2003). Reprinted with permission from Tits et al. (2003). Copyright 2005 by Walter de Gruyter GmbH.

The incorporation processes observed in the spectroscopic studies appear to have no significant influence on the uptake kinetics and on the desorption behaviour. No sorption hysteresis was observed. Eu(III)/Cm(III) uptake into the C–S–H interlayers might be a fast and reversible process comparable to ion exchange processes in clay interlayers. Structural incorporation into the Ca-O layers, however, is expected to be a slow process with uptake rates comparable to the recrystallization rates measured by 45Ca exchange. The 45Ca recrystallization experiments showed that after an equilibration period of 60 days, approximately 30% of the C–S–H material was recrystallized in ACW (Fig. 5c). The Eu(III) uptake kinetic experiments, however, do not provide any evidence for such slow kinetics. Maybe this slow uptake is hidden by the large uncertainties on the measured Rd values. Assuming that all Ca sites in the C–S–H phases can be occupied by trivalent actinides and that the affinity of An(III)/Ln(III) and Ca2+ for these sites is identical, it was estimated that the amount of An(III) in solution after 60 days would be reduced by a factor 4 due to recrystallization (Fig. 5b). Such a reduction would very likely not be detected within the large scatter in the sorption data.

580 590 600 610 620 630 640

Contact time

119 d 36 d 15 d 10 d 4 d 2 d 20 h 1 h

0.05 g L-1 CSH1 x 10-7 M Cm(III)

Cm(II

I) flu

ores

cenc

e em

issio

n

Wavelength (nm)580 590 600 610 620 630 640

shift

Norm

alize

d flu

ores

cenc

e em

issio

n

Wavelength (nm)


To conclude, trivalent actinides and lanthanides may be incorporated into the Ca-O main layers of C–S–H phases but the effects of this incorporation process are most probably hidden by the large uncertainties on the experimental data. A conclusive understanding of the detailed incorporation mechanism is still missing.

6 Sorption of tetravalent actinides

6.1 Speciation

The main tetravalent actinides relevant for radioactive waste are U(IV), Np(IV) and Pu(IV). In addition, Th(IV) is often used as a chemical analogue for the two former tetravalent actinides because it is the only actinide having a stable +IV redox state under oxidizing conditions thus facilitating significantly experimental work. Comparison of the complexation constants of Th(IV) and other tetravalent actinides (e.g. Np(IV) in Table 3), however, shows that this analogy can be questioned. Nonetheless, the majority of the sorption experiments with tetravalent actinides on cementitious materials discussed in the present report were carried out with Th(IV) because of its redox stability. In some cases Th(IV) sorption experiments were complemented with Np(IV) sorption experiments in the presence of Na-dithionite to stabilize the +IV redox state. The following speciation calculations were carried out with Th(IV) in cementitious pore waters at varying compositions using the code “Medusa” (Puigdomenech, 1983). The case of Th(IV) was taken because only for this tetravalent actinide stability constants for silicate complexes were available and it is shown that these complexes play an important role in the speciation of these actinides together with the tetra hydroxo species. The stability constants for the Th(IV) hydroxy complexes were taken from the NEA database of Rand et al. (2008). The stability constants for the ternary Ca-Th(IV)-OH complex and the ternary Th(OH)3(SiO(OH)3)3

2- complex were taken from Altmaier et al. (2008) and Rai et al. (2008), respectively. The solubility of Th(IV) in alkaline conditions was calculated by taking into account solely the solubility product of the fresh amorphous ThO2 phase,ThO2(am, fresh, hyd) (Guillaumont et al., 2003). The aged amorphous ThO2 phase, ThO2(am, aged, hyd), as described in the NEA thermodynamic database for Thorium (Rand et al., 2008) was neglected.. All constants involving reactions with Th(IV) are listed in Table 3. For comparison, the stability constants for the equivalent Np(IV) complexes are listed in Table 3 as well (Guillaumont et al., 2003; Fellhauer et al., 2010). Constants for all other reactions needed in the calculations were taken from the core dataset in Hummel et al. (2002). Extrapolation to I = 0.3 M was done using the Davies equation (b = 0.3) (Davies, 1962).

Speciation calculations were carried out for three different pore waters to determine the maximal effect of Ca and Si(OH)4 on the An(IV) speciation: 1) A pore water with high Ca concentration (0.02 M) and low Si concentration (10-4 M), reflecting a cement pore water in equilibrium with a C–S–H phase with high C:S ratio, 2) a pore water with lower Ca concentration (10-3 M) and intermediate Si concentration (10-3 M) and 3) a pore water with a high Si(OH)4 concentration (8·10-3 M) and low Ca concentration (4·10-5 M), reflecting a pore water in equilibrium with a C–S–H phase with low C:S ratio. All calculations were performed by fixing I to 0.3 M in the code Medusa. This value reflects the ionic strength in ACW. Note that this implies that the speciation calculated at pH > 13.3 is only approximately correct.

The calculated concentrations of dissolved Th(IV) are shown in Fig. 14. Aqueous silicate species and aqueous Ca species are omitted to simplify the presentation of data. At the highest Si concentrations, the Th(OH)3(SiO(OH)3)3

2- species dominates the speciation in the pH range


10.0 < pH < 13.0 and the Th(IV) solubility increases with decreasing pH to reach a value of approximately 10-6 M at pH = 10.0. At pH values above 13.0, the hydrolysed Th(OH)4(aq) species replaces Th(OH)3(SiO(OH)3)3

2- as the dominant species. At intermediate Si concentrations below 10-3 M, the influence of the Th(OH)3(SiO(OH)3)3

2- species on the speciation decreases drastically, and at low Si concentrations of 10-4 M, the dominating aqueous Th(IV) species is Th(OH)4(aq) and the solubility of ThO2(am, fresh, hyd) is predicted to be approximately 10-8 M independent of pH. The ternary Ca-Th(IV)-OH complex does not play a role in the Th(IV) speciation in cementitious pore waters. Thus, it is concluded that in fresh ACWHCP-I (pH = 13.3) and in degraded ACWHCP-II (pH = 12.5) hydrolysed Th(OH)4 species will dominate the aqueous Th(IV) speciation. In low pH cements containing large amounts of silica fume, however, the high aqueous Si concentrations could significantly increase the ThO2(am, fresh, hyd) solubility.

Table 3 Equilibrium constants for An(IV) aqueous hydroxide complexes, ternary Ca-An(IV)-OH complexes and solid An(IV) hydroxides (Guillaumont et al., 2003; Altmaier et al., 2008; Rai et al., 2008; Rand et al., 2008; Fellhauer et al., 2010; Thoenen et al., 2014).

Hydroxide complexes: log*β0n:

4 4 n2 nAn n H O An(OH) n H+ − ++ ⋅ ⇔ + ⋅

Th(IV) Np(IV)

Th(OH)3+ -10.85

Th(OH)4(aq) -17.4 Np(OH)4(aq) -8.3

Ternary Ca-An(IV)-OH complexes: log*β0p, 1, n:

2 4 2p 4 n2 p npCa An nH O Ca [An(OH) ] nH+ + + − ++ + ⇔ +

44 8Ca [Th(OH) ] + -62.7 4

4 8Ca [Np(OH) ] + -56.1

Ternary An(IV)-OH-Si complexes: log01,p,n :

4 n p44 2 n 3 pAn pSi(OH) (aq) nH O Th(OH) (SiO(OH) ) 2nH− −+ ++ + ⇔ +

Th(OH)3(SiO(OH)3)32- -27.8

Solubility: log*K0s:

42 2An 4OH AnO (am,hyd) 2H O+ −+ ⇔ +

ThO2(am, fresh, hyd)

ThO2(am, aged, hyd)

46.7

47.5

NpO2(am, hyd) 56.7


Fig. 14: Tetravalent actinide speciation as a function of pH in the absence of carbonates. [Th(IV)]tot = 10-5 M. I = 0.3 M. a) [Ca] = 5·10-5 M and [Si(OH)4] = 8·10-3 M, b) [Ca] = 10-3 M and [Si(OH)4] = 10-3 M, c) [Ca] = 2·10-2 M and [Si(OH)4] = 10-4 M. Thermodynamic equilibrium constants for solid hydroxides and aqueous hydroxyl complexes are taken from the NEA thermodynamic databases (Guillaumont et al., 2003; Rand et al., 2008) and from Fellhauer et al., (2010).

6.2 Preliminary solubility tests

The results of Th(IV) solubility tests carried out in ACW are displayed in Figs. 15a and b. Note that the ACW used in these solubility tests does not contain Si because actinides – silicate complexes were considered to be irrelevant at the time these experiments were carried out. The figures show the ratio of the radionuclide concentration determined in solution to the total added radionuclide concentration (Fig. 15a) and the solution concentrations of the radionuclide as a function of the added radionuclide concentration in solution (Fig. 15b). The data refer to measured concentrations before centrifugation (1 h at 95,000g) and after centrifugation. Particles with diameter larger than ~3 nm are expected to settle during centrifugation. Thus, the concentration determined after centrifugation represents the concentration of dissolved radionuclides plus the concentration of radionuclides associated with colloidal material smaller than ~3 nm.

The reduction of the Th(IV) concentration in solution before centrifugation was found to range in value between 20% and 80% of the initial Th(IV) concentration added to ACW. This reduction was observed over the entire concentration range and is caused by wall sorption. The

10 11 12 13 14-10

-9

-8

-7

-6

-5

-4

Th(IV) solubility

Log

conc

entra

tion

(M)

pH

Th(OH)4(aq)

[Ca]=0.02 M[Si(OH)4]=10-4 M

c


percentage of the solution concentration before centrifugation was found to be consistently lower after 21 days equilibration than after 1 day equilibration. This indicates that Th(IV) sorption onto the container walls is a slow process. Centrifugation caused an additional reduction in solution concentration at the highest Th(IV) concentrations due to the precipitation of Th(IV) colloids.

In Fig. 15b the continuous line with slope 1 represents the case where the initial concentration in solution is completely recovered in solution. The data obtained before centrifugation closely follow this line. At starting Th(IV) concentrations ≥ 10-8 M, however, a constant Th(IV) concentration in solution of ~6·10-9 M was determined after centrifugation (solution stability limit). This can be explained by the formation of Th(IV) radiocolloids or a Th(IV) solid phase, respectively, which are formed in ACW and precipitate during centrifugation. The observed solution concentration is very similar to the Th(IV) solubility w.r.t. ThO2(am, fresh, hyd) in ACW (shaded area in Fig. 15b). Based on these measurements it was concluded that the initial Th(IV) concentration added to ACW should be below 10-8 M to avoid the formation of colloidal ThO2(s). Solubility tests in alkali-free solutions at lower pH were not carried out. Nevertheless, the same solubility limit is expected in the pH range 10 to 13.3 due to identical speciation (Fig. 14c). Furthermore, the stability of Np(IV) in alkaline conditions is assumed to be similar to the Th(IV) stability.

Fig. 15: Solubility tests with Th(IV) in ACW at pH = 13.3. a) Percentage of the Th(IV) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Th(IV) inventory. b) Th(IV) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Th(IV) inventory. Equilibration times were 1 day and 21 days. The curves in a) only serve to guide the eye. The shaded area in b) represents the calculated Th(IV) solubility limit with its uncertainty in ACW with respect to ThO2(am, fresh, hyd). The line with slope +1 represents 100% recovery in solution.

6.3 Sorption and co-precipitation kinetics

6.3.1 Sorption kinetics The results from Th(IV) and Np(IV) sorption kinetic tests in ACW (pH = 13.3) and in alkali-free solutions (pH = 10.0 - 12.5) with C–S–H phases are shown in Figs. 16a – 16d. Note that, for all except one experimental series, the Th(IV) and Np(IV) concentrations used were well

10-11 10-10 10-9 10-8 10-7 10-6 10-5

0

20

40

60

80

100

1 day; before centr. 1 day; after centr. 21 days; before centr. 21 days; after centr.

% o

f add

ed T

h(IV

) mea

sure

d in

equi

libriu

m so

lutio

n

Concentration of added Th(IV) (M)

a

10-11 10-10 10-9 10-8 10-7 10-6 10-510-11

10-10

10-9

10-8

10-7

10-6

10-5

1 day before centr. 1 day after centr. 21 days before centr. 21 days after centr.

ThO2(am, fresh, hyd)

Th(IV

) con

cent

ratio

n m

easu

red

in so

lutio

n (M

)

Concentration of added Th(IV) (M)

b

Slope = +1


below the solubility limits w.r.t. the respective amorphous hydrated oxides as well as below the stability limits measured in the stability experiments. Note that formation of amorphous hydrated oxides was not observed even in the only experiment where a total Np(IV) concentration above the calculated solubility limit was used. Furthermore, a large uncertainty on the measured distribution ratios was observed for the strongly sorbing Th(IV) and Np(IV) which is attributed to incomplete phase separation. The Rd,max values were estimated to be 1.16·107 L kg-1 and 2.8·108 L kg-1 for the set-up in the experiments with 228Th(IV) and 239Np(IV), respectively.

Fig. 16: Np(IV) and Th(IV) sorption kinetics on C–S–H phases having various compositions in alkali-free conditions and in ACW at pH = 13.3. a) Np(IV) sorption in alkali-free conditions (pH 10.1), b) Np(IV) sorption in ACW (pH 13.3), c) Th(IV) sorption in alkali-free conditions (pH 10.1), d) Th(IV) sorption in ACW (pH 13.3). Experimental conditions: Np(IV) sorption kinetics: S:L = 2·10-4 kg L-1, [Np(IV)]tot = 1.7·10-7 M (237Np + 239Np) and 1.3·10-10 M (239Np). Th(IV) sorption kinetics: S:L = 2·10-3 kg L-1, [Th(IV)]tot = 2·10-11 M.

The uptake behaviour of Th(IV) and Np(IV) was very similar under all experimental conditions. The Rd values were very high (Rd > 105 L kg-1) both in ACW (Figs. 16a and b) and in the pH range between 10 and 12.5 (Figs. 16c and d). Furthermore, uptake was fast, and equilibrium was reached within one day. The Rd values were constant within the uncertainty over the investigated time period (up to 120 days for Th(IV) and up to 10 days for Np(IV)). The Rd

2 4 6 8 10104

105

106

107

108 C:S=1.65; [Np(IV)]tot = 1.7·10-7 M

C:S=1.65; [Np(IV)]tot = 1.3·10-10 M

C:S=0.65; [Np(IV)]tot = 1.3·10-10 M

R d(N

p(IV

)) (L

kg-1

)

Time (days)

Alkali-free conditions

aRd,max

0 2 4 6 8 10104

105

106

107

108

C:S = 0.65 C:S = 1.65

R d(N

p(IV

)) (L

kg-1

)

Time (days)

ACW

bRd,max

0 20 40 60 80 100 120104

105

106

107

108

C:S=0.75; pH=10.4 C:S=1.08; pH=12.1 C:S=1.28; pH=12.3 C:S=1.6; pH=12.5

R d (T

h(IV

)) (L

kg-1

)

Time (days)


c

Rd,max

0 20 40 60 80 100 120104

105

106

107

108

C:S=0.75 C:S=1.08

R d (T

h(IV

)) (L

kg-1

)

Time (days)

ACW

d

Rd,max


values are considered to be independent of the C:S ratios in view of the uncertainties on the data. The formation of aqueous An(IV) – silicate complexes that could occur according to the speciation calculations (5·10-6 M < [Si] < 8·10-3 M) appears to have no effect on the An(IV) sorption in the experiments with C–S–H phases having low C:S ratios at pH = 10.1 in alkali-free conditions (see Fig. 14a and 14c). In ACW (pH value = 13.3) the concentration of aqueous ternary An(IV)-OH-Si(OH)4 complexes is negligible while the aqueous Th(IV) speciation is dominated by the hydrolysed Th(OH)4 species (Fig. 14c).

The Th(IV) sorption kinetics onto HCP resembles that on synthetic C–S–H phases (Fig. 17): The measured Rd values were found to be a factor 2 to 5 lower than those determined on C–S–H phases. This lower value may be related to the HCP composition which consists of only ~50% C–S–H phases, the main sorbing phase in HCP. Rd values for aged HCP (pH = 12.5) were slightly higher than Rd values for fresh HCP at pH = 13.3. Nevertheless, this difference is hardly significant in view of the uncertainty in the sorption measurements.

Fig. 17: Th(IV) sorption kinetics onto fresh HCP in ACWHCP-I (pH = 13.3) and onto aged HCP (ACWHCP-II, pH = 12.5). Experimental conditions: S:L = 1.0·10-4 kg L-1, [Th(IV)]tot = 2·10-10 M.

6.3.2 Th(IV) co-precipitation kinetics In addition to sorption kinetic tests, a series of co-precipitation kinetics tests were carried out with Th(IV) on C–S–H phases with varying C:S ratios in alkali-free solutions and in ACW (Figs 18a and b). Rd,max values were calculated to be similar to those in the sorption kinetic experiments (Section 6.3.1.). It was observed that the Th(IV) co-precipitation kinetics is very similar to the sorption kinetics, i.e. fast Th(IV) uptake reaching steady state within 1 day. The Rd values obtained from co-precipitation experiments, however, tend to be slightly higher (~106 L kg-1 < Rd < ~107 L kg-1) than those determined in the sorption experiments (~105 L kg-1 < Rd < ~106 L kg-1) (see Figs. 16 and 18). Furthermore, the Rd values determined in the co-precipitation experiments are independent of the C:S ratios in view of the uncertainties on the data. Note that several experimental data in Figs. 18a and b, are higher than the Rd,max value calculated for these experiments. These data are above the maximal Rd value, statistically measurable with the present experimental set-up and should therefore be interpreted with caution.

0 20 40 60 80 100 120 140104

105

106

107

ACWeq-I pH = 13.3; exp. 1 ACWeq-I pH = 13.3; exp. 2 ACWeq-II pH=12.5

R d(Th(

IV))

(L k

g-1)

Time (days)


Fig. 18: Th(IV) co-precipitation kinetics onto C–S–H phases with various C:S ratios in a) alkali-free systems and b) ACW (pH 13.3). Experimental conditions: S:L = 2·10-3 kg L-1, [Th]tot = 2·10-11 M. Comparison with the data from the sorption kinetic experiments (datapoints in grey)

6.4 Desorption experiments

Desorption experiments with Th(IV) were carried out on a C–S–H phase with a C:S ratio of 1.07 (Figs. 19a and b) in ACW at pH = 13.3. Desorption was started after 1 day sorption, 20 days sorption and 120 days sorption. In each case three subsequent desorption steps were carried out. The equilibration time during desorption was 3 days. Note that Rd,max values are similar to those in the sorption kinetic experiments (Section 6.3.1.).

Fig. 19: Th(IV) sorption (a) and co-precipitation (b) experiments on a C–S–H phase with a C:S ratio of 1.07 in ACW at pH 13.3. Desorption tests conducted after 1 day, 20 day and 120 days of sorption/co-precipitation. Experimental conditions: S:L = 2·10-3 kg L-1, [Th]tot = 2·10-12 M, desorption time = 3 days.

The Rd values obtained from desorption tests were found to be slightly higher than those determined in the sorption and in the co-precipitation experiments. Note that the interpretation of the results of the desorption experiments is difficult as the aqueous Th(IV) concentrations in the supernatant solutions were extremely low so that minimal variations in the measured concentrations (e.g. due to incomplete phase separation) had a strong effect on the Rd values. The slightly higher desorption Rd values observed in these experiments should therefore be

0 20 40 60 80 100 120103

104

105

106

107

108

C:S=0.75; pH=10.4 C:S=1.08; pH=12.1 C:S=1.28; pH=12.3 C:S=1.65; pH=12.5

R d (T

h(IV

)) (L

kg-1

)

Time (days)

Alkali-free conditionsaRd,max

0 20 40 60 80 100 120104

105

106

107

108

C:S=0.75 C:S=1.08

R d (T

h(IV

)) (L

kg-1

)

Time (days)

ACWb

Rd,max

0 20 40 60 80 100 120 140104

105

106

107

108

Rd (sorption) Rd (1

th desorption) Rd (2


th desorption)

R d (T

h(IV

)) (L

kg-1

)

Time (days)

C:S = 1.08 in ACW

a

Rd,max

0 20 40 60 80 100 120 140 160104

105

106

107

108

C:S = 1.08 in ACW

Rd (co-precipitation) Rd (1



th desorption)

R d (T

h(IV

)) (L

kg-1

)

Time (days)

bRd,max


interpreted with caution. With a view to the observed absence of any hysteresis in the sorption of the trivalent actinides reported in section 5.4., it is anticipated that the Th(IV) sorption observed in this study is reversible as well.

6.5 Effect of S:L ratio

The sorption studies with Eu(III) showed that the measured Rd values may depend on the S:L ratio of the experimental system. In the case of Th(IV) we checked Th(IV) sorption on C–S–H phases having three different C:S ratios (Fig. 20a) and on fresh HCP (Fig. 20b) in suspensions with different S:L ratios. Rd,max values for Th(IV) sorption onto C–S–H phases with C:S ratios 0.96 and 1.29 are represented by a black dashed line whereas the Rd,max values for Th(IV) sorption onto C–S–H phases with a C:S ratio of 1.25 is represented by the red dash-dotted line. The Rd,max value of the latter experiments is much higher because the 228Th analysis in solution was carried out using a liquid scintillation counter equipped with α/β discrimination resulting in a much lower background (3 cpm/5 mL compared to 200 cpm/5 mL).

The effect of the S:L ratio of the C–S–H suspensions on the sorption of Th(IV) was negligibly small. In the case of fresh HCP, a slight trend towards decreasing Rd values with increasing S:L ratios is suggested. We have presently no explanation for this behaviour in HCP. In view of the low reproducibility of the experimental sorption data, and in view of the fact that no effect of the S:L ratio was observed in the experimental systems with C–S–H phases, we infer that there is no major effect of the S:L ratio on An(IV) sorption.

Fig. 20: Th(IV) sorption as a function of the S:L ratio in ACW at pH = 13.3 on C–S–H phases with different C:S ratios (a) and on fresh HCP (b). Experimental conditions for the C–S–H systems: [Th]tot = 5·10-10 M, equilibration time = 2 weeks. Experimental conditions for fresh HCP in ACWHCP-I: [Th]tot = 7·10-10 M, equilibration time = 2 weeks.


A Th(IV) sorption isotherm was determined for a C–S–H phase with C:S ratio of 1.07 in ACW at pH 13.3 (Figs. 21a and b). Rd,max values are shown as dashed lines in Fig. 21b. Equilibrium Th(IV) concentrations varied between 6·10-14 M and 10-10 M. Comparison of the measured Rd values of the isotherm with the Rd,max values shows that some of the experimental data were very close to the experimentally measurable values.

10-5 10-4 10-3 10-2 10-1103

104

105

106

107

C:S = 0.96 C:S = 1.25 C:S = 1.29

Rd (

Th(I

V))

(L k

g-1)

S:L (kg L-1)

a

10-5 10-4 10-3 10-2 10-1103

104

105

106

107

Exp. 1 Exp. 2 Exp. 3

Rd (

Th(I

V))

(L k

g-1)

S:L (kg/L)

b


This isotherm was found to be linear over the entire concentration range investigated. The highest Th(IV) loading covered by the sorption isotherm is 2·10-5 mol kg-1 or 8·10-5 eq kg-1 (assuming that one Th(IV) neutralizes four sorption sites on the C–S–H phase). This Th(IV) loading corresponds to ~0.001% of the sorption capacity of the C–S–H phases (~6.2 eq kg-1; see Section 2.1.2). This low loading implies that only the sorption sites with the highest affinity for Th(IV) are occupied, which explains the linear sorption behaviour of Th(IV).

Th(IV) sorption isotherm measured on fresh HCP at pH = 13.3 (Fig. 21c) and on degraded HCP at pH = 12.5 (Fig. 21d) showed a very similar behaviour; i.e. a linear sorption behaviour over a [Th]eq concentration range between 10-14 M and 10-8 M and no influence of the degradation stage (pH= 13.3 versus pH = 12.5). In the degraded HCP system at pH 12.5, precipitation of ThO2(am, fresh, hyd) appears to start at slightly lower equilibrium Th(IV) concentrations than in fresh HCP at pH 13.3. Nevertheless, this difference is considered insignificant.

Fig. 21: a, b) Th(IV) sorption isotherm onto a C–S–H phase with C:S = 1.07. Experimental conditions: S:L = 5·10-4 kg L-1, equilibration time = 2 weeks. c, d) Th(IV) sorption isotherm onto fresh HCP in ACWHCP-1 (pH = 13.3) and in aged HCP in ACWHCP–II (portlandite saturated solution, pH = 12.5). Experimental conditions: S:L = 10-4 kg L-1, equilibration time = 2 weeks. Black vertical lines: Experimental Th(IV) stability limit. Red vertical lines: Thermodynamic Th(IV) solubility limit.

10-14 10-13 10-12 10-11 10-10 10-9 10-810-9

10-8

10-7

10-6

10-5

10-4C:S = 1.07 in ACW

Th(IV

) sorb

ed (

mol

es k

g-1 )

Th(IV) equilibrium concentration (M)

Slope = 1.0

a

10-14 10-13 10-12 10-11 10-10 10-9 10-8104

105

106

107

R d(T

h(IV

)) (L

kg-1

)


Mean Rd = 3±2x105 L kg-1

C:S = 1.07 in ACW

Th(IV

) stab

ility

lim

it

b

Rd,max

10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-610-9

10-8

10-7

10-6

10-5

10-4

10-3

10-2

10-1

100


ACWHCP-I pH=13.3 ACWHCP-II; pH=12.5

Th(IV

) sorb

ed (

mol

es k

g-1 ) c

10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6104

105

106

107

108ACWHCP-I; pH = 13.3 ACWHCP-II; pH = 12.5

R d(T

h(IV

)) (L

kg-1

)


Mean Rd = (3±2)x105 L kg-1

d


6.7 Effect of the C–S–H composition (C:S ratio)

In order to assess the effect of the aqueous speciation on the sorption of Th(IV) and Np(IV) cations, the uptake of Th(IV) and Np(IV) by C–S–H phases was determined as function of the C:S ratio under alkali-free conditions and in ACW at pH = 13.3 (Fig. 22). The results show that the Rd values appear to be independent of the C–S–H composition. The C:S ratio of C–S–H phases in alkali-free conditions is positively correlated with both the pH and the Ca concentration in solution and inversely correlated with the Si concentration (see Section 4.2). We assume that the variation in the C–S–H structure has no influence upon the Th(IV) sorption as the total number of sorption sites is almost independent of the C:S ratio (see Section 2.1.2.).

The sorption of Th(IV)/Np(IV) is independent of pH. In alkali-free conditions the Rd value remains constant when the pH rises from 10.1 to 12.5. This observation is in line with the idea that the same aqueous species are involved in the partitioning of Th(IV)/Np(IV) between C–S–H phases and solution and therefore, further in line with the speciation calculations that show Th(OH)4(aq)/Np(OH)4(aq) as the dominant aqueous species.

The increase in Ca2+ concentration from ∼3·10−5 M to 1.6·10−3 at constant pH of 13.3 in ACW (due to the increase in C:S ratio) obviously does not exert an effect on the Th(IV) Rd value indicating that the uptake of tetravalent actinides is not accompanied by the release of Ca2+ and that the uptake can most likely not be explained by a simple Ca2+ – Th(IV)/Np(IV) ion exchange process in the C–S–H interlayer.

Fig. 22: An(IV) sorption onto C–S–H phases as function of the C:S ratio in alkali-free solutions and in ACW (pH 13.3). Experimental conditions for Np(IV) experiments: [Np]tot = 2·10-10 M, S:L = 5·10-3 kg L-1, equilibration time = 3 days. Experimental conditions for the Th(IV) experiments: [Th]tot = 8.5·10-9 M, S:L = 5·10-3 kg L-1, equilibration time = 3 days.

6.8 Discussion

6.8.1 Sorption data The sorption behaviour of An(IV), in particular Th(IV) and Np(IV), on C–S–H and HCP can be summarized as follows:

0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0102

103

104

105

106

107

Np(IV) in H2OTh(IV) in H2O Np(IV) in ACW Th(IV) in ACW

R d (L k

g-1)

C:S (mol/mol)


1) Very strong sorption illustrated by Rd values ranging between 105 L kg-1 < Rd < 3·106 L kg-1;

2) Co-precipitation gives rise to slightly higher Rd values than sorption (106 L kg-1 < Rd < 107 L kg-1);

3) Fast sorption and co-precipitation kinetics without sorption hysteresis onto C–S–H phases and HCP;

4) The sorption process is linear in the concentration range investigated;

5) The uptake process of tetravalent actinides by C–S–H phases is independent of the C:S ratio;

6) No effect of pH on the sorption of tetravalent actinides in the pH range 10 – 13.3 due to constant aqueous speciation;

7) No effect of cement degradation on tetravalent actinide sorption as expected from the previous observations.

The reviews of Ochs et al. (2016) and Wieland (2014) provide a comprehensive overview of the research on the An(IV) sorption onto cementitious materials available in the open literature. Earlier studies discussed in these reviews were usually carried out using higher S:L ratios (5·10-3 kg L-1 to 5·10-2 kg L-1) and report lower Rd values (~103 L kg-1 – 104 L kg-1). In more recent studies (e.g. Pointeau et al., 2004a), it was recognised that application of low S:L ratios is mandatory in sorption experiments with strongly sorbing tracers, to obtain reliable Rd values. In these recent studies higher Rd values were reported in agreement with the Rd values found in the present study. Effects of the S:L ratio (decreasing Rd with increasing S:L ratio) and of the cement degradation (increasing Rd values with progressing degradation) previously claimed by Ochs et al., (2016) are not conclusively supported by our data.

6.8.2 Uptake mechanisms The absence of any effect of pH on sorption is in line with the finding that the fourfold hydrolysed species, Th(OH)4 or Np(OH)4, are the dominant species in solution and thus the main sorbing species. The theory of the electrostatic inter ligand repulsion (Section 4) suggests that the limiting number (nlimit) of OH groups that fit in the first coordination sphere of An(IV) is larger than 4. Thus, Th(OH)4 and Np(OH)4 are able to form strong bonds with the silandiol and silanol sites in the interlayers and at the edges of the C–S–H phases, which supports the assumption that the fourfold hydrolysed species is the sorbing species.

Recent X-ray Absorption Fine-Structure spectroscopy (XAFS) investigations provided ample evidence for the incorporation of Np(IV) in the interlayer of C–S–H phases (Gaona et al., 2011). The authors observed Np(IV) incorporation into the C–S–H structure in two different coordination environments depending on the C:S ratio of the C–S–H phases. While at low C:S ratios, higher Si coordination numbers and lower Ca coordination numbers were observed, the Si coordination numbers decreased and the Ca coordination numbers increased at higher C:S ratios, consistent with intercalation of Np(IV) in the C–S–H interlayers. The batch sorption data, however, suggest that another process rather than a simple Ca – Np(IV) ion exchange process in the C–S–H interlayers controls the Np(IV) uptake. A variation in the C:S ratio from C:S = 0.65 to 1.60, at constant pH of 13.3 in ACW, implying a variation of equilibrium Ca concentration over more than two orders of magnitude, had no measurable effect on the Rd value. This observation suggests that the Np(IV) species and Ca2+ occupy different crystallographic sites in the interlayer, at least at those sites where Ca2+ bound in the structure is in fast equilibrium with aqueous Ca.


The sorption kinetics and the desorption studies finally show that tetravalent actinide uptake by the C–S–H interlayers has no influence on the kinetics and that it is a fully reversible process. Overall, Th(IV) and Np(IV) incorporation into the C–S–H interlayer, clearly indicated by spectroscopy, appears to have no measurable effect on their macroscopic sorption behaviour.

7 Sorption of pentavalent actinides

7.1 Speciation

Actinides having stable pentavalent oxidation states are mainly Np and Pu, which form linear dioxo moieties (O=An=O+), and Pa having a speciation strongly different from Np and Pu. The present work only focusses on the sorption of Np(V). Pu(V) is expected to exhibit a very similar sorption behaviour. The sorption of Pa(V) was not investigated.

Speciation calculations were performed with the code “Medusa” (Puigdomenech, 1983). The thermodynamic complexation constants were taken from the NEA database (Guillaumont et al., 2003) and completed with the thermodynamic data for ternary Ca-NpO2-OH complexes reported by Fellhauer et al. (2016a) and with the thermodynamic data for NpO2-silicate complexes reviewed by Thoenen et al. (2014). All these data are listed in Table 4. Thermodynamic complexation constants for all other reactions needed in the calculations were taken from the core dataset reported in Hummel et al. (2002). All calculations were performed fixing I to 0.3 M in the code Medusa. This value reflects the ionic strength in ACW. Note that this implies that the speciation calculated at pH > 13.3 is only approximately correct. The speciation of Np(V) in alkaline conditions between pH 10.0 and 13.3, in the absence of Ca and Si, is determined by the hydrolysis and dominated by three species: NpO2

+, NpO2OH(aq) and NpO2(OH)2

- (Figs. 23a). The solubility controlling Np(V) phase is NpO2OH(am, fresh) (Guillaumont et al., 2003). Recent investigations confirmed this thermodynamic model as no evidence for the ageing of the latter solid phase and its transformation into NpO2OH(am, aged) could be provided (Petrov et al., 2017). Hence, the latter solid was not considered for the speciation calculations. The same authors provided evidence for the existence of ternary Na-Np(V)-OH phases in concentrated NaCl systems (>1 M NaCl at pH > 11). These phases are, however, not relevant in cementitious environments and are ignored in this study.

In the presence of Ca, the speciation of Np(V) is completely different than in the absence of Ca (Altmaier et al., 2013; Fellhauer et al., 2016a; 2016b). Figs. 23b and c show the results of Np(V) solubility and speciation calculations in alkaline conditions at two different Ca concentrations, 4·10-5 M Ca and 0.02 M Ca, i.e. the range of Ca concentrations expected in solutions in equilibrium with C–S–H phases with C:S ratios of 0.66 and 1.8, respectively, and further relevant to HCP systems. Note that the total Np(V) concentration in these calculations was fixed at a value below the total Ca concentration (realistic assumption) but still high enough to allow precipitation of solid Ca-Np(V) phases. Already at the lowest Ca concentration (Fig. 23b), the Np(V) solubility decreases significantly compared to the system without Ca (Fig. 23a) due to the formation of a solid ternary Ca-Np(V)-OH compound: Ca0.5NpO2(OH)2·1.3H2O(s). However, the aqueous Np(V) speciation under these chemical conditions is still dominated by the same Np(V) hydroxy complexes as in the system without Ca (NpO2OH and NpO2(OH)2

-). At higher Ca concentrations, the Np(V) solubility is even lower and estimated to be ~7·10-10 M at pH = 12.5. Furthermore, a ternary Ca-Np(V)-OH complex, Ca[NpO2(OH)2]+ is stable in the aqueous phase. The presence of Ca in cement pore water thus appears to reduce significantly the solubility of Np(V) in cementitious environments.


Table 4: Summary of log10*β° and log10*K°s,0 used in the Np(V) speciation calculations.

Only species formed between 10 < pH < 14 are considered. (Guillaumont et al., 2003; Altmaier et al., 2013; Thoenen et al., 2014; Fellhauer et al., 2016a)

Species log10*βn° Reaction

NpO2OH(aq)

NpO2(OH)2-

Ca[NpO2(OH)2]+

Ca3[NpO2(OH)5]2+

NpO2SiO(OH)3

-11.3

-23.6

-20.63

-54.81

7.04

2 22NpO H O NpO OH(aq) H+ ++ ⇔ +

2 2 22NpO 2H O NpO (OH) 2H+ − ++ ⇔ +

22 22 2Ca NpO 2H O Ca [NpO (OH) ] 2H+ + + ++ + ⇔ +

2 22 3 2 523Ca NpO 5H O Ca [NpO (OH) ] 5H+ + + ++ + ⇔ +

3 22 3NpO SiO(OH) NpO SiO(OH) (aq)+ −+ ⇔ -

Solid Log10*K°s,0 Reaction

NpO2OH(am, fresh)

Ca0.5NpO2(OH)2·1.3H2O(s)

-5.3

-12.3

2 22NpO H O NpO OH(am,fresh) H+ ++ ⇔ +

22 0.5 2 22 20.5Ca NpO 3.3H O Ca NpO (OH) 1.3H O(s) 2H+ + ++ + ⇔ ⋅ +

Fig. 23: Np(V) speciation calculations in the pH range 10 < pH < 14 in the absence of Ca (a) and in the presence of low (b) and high (c) Ca concentrations representing the typical Ca concentrations in C–S–H solutions in the absence of alkalis.

10 11 12 13 14-8

-7

-6

-5

-4a

NpO2(OH)-2

NpO2OH(aq)

Log[

Np(

V)]

(M)

pH

[NpO+2]tot=10-5 M

[Ca]tot = 0.0 M

NpO2OH(am, fresh)NpO+2

10 11 12 13 14-12

-10

-8

-6

-4

-2


Ca[NpO2(OH)2]+

NpO2(OH)-2

CaOH+

NpO2OH

NpO+2

Log[

Np(

V)]

(M)

pH

[Np(V)]tot = 10-5M[Ca]tot = 5·10-5 M

b

Ca2+

10 11 12 13 14-10

-9

-8

-7

-6

Ca3[NpO2(OH)5]+2

Ca(OH)2 precipitation

c

NpO2(OH)-2

Ca[NpO2(OH)2]+

NpO2OH(aq)

Log

[Np(

V)]

pH

[NpO+2]tot=10-5 M

[Ca]tot = 0.02 M


NpO+2


The An(V)-silicate complexes that predominantly exist in the low pH region (e.g. NpO2SiO(OH)3(aq)) do not play a role in the Np(V) speciation distribution in alkaline conditions. The existence of ternary Np(V)–hydroxy–silicate complexes expected to prevail in the alkaline pH region has not yet been demonstrated. These complexes could play some role in the pore waters of low pH cements where the aqueous Si concentration is high.


Np(V) solubility tests were carried out in alkali-free solutions containing 10-4 M Ca(OH)2 at pH = 10.3 in 0.3 M NaOH at pH = 13.3 and in ACW at pH = 13.3. The solutions were equilibrated on an end-over-end shaker for 30 days prior to sampling. The figures show the measured Np(V) solution concentrations before and after centrifugation (1 h at 95,000g) as a function of the radionuclide concentration added to the solution (Figs. 24a,b and c). Note that only particles with diameter larger than ~3 nm are expected to settle during centrifugation (see Section 3.3.3). Thus, the concentration determined after centrifugation represents the concentration of dissolved radionuclides plus the concentration of radionuclides associated with colloidal material smaller than ~3 nm.

In the alkali-free solutions containing 10-4 M Ca(OH)2 at pH = 10.3, the Np(V) concentrations measured before and after centrifugation are identical to the Np(V) concentrations added to the solutions over the entire investigated concentration range (Fig. 24a). The Np(V) solubility in this system is controlled by Ca0.5NpO2(OH)2·1.3H2O and was calculated to be 2.5·10-6 M, i.e. a factor of 4 below the highest Np(V) concentration tested. The absence of a Ca0.5NpO2(OH)2·1.3H2O precipitate at the highest Np(V) concentration (10-5 M) is assumed to be due to incomplete phase separation during the centrifugation step.

In the system with higher Ca concentrations (0.01 M Ca(OH)2, pH 12.3), the Np(V) concentrations measured in the solutions before and after centrifugation were significantly lower than the concentrations originally added to the solutions (Fig. 24b). The lower Np(V) concentration before centrifugation is most probably caused by Np(V) sorption on the walls of the centrifuge tubes. This assumption is further supported by the observation that the missing Np(V) could be recovered after washing the centrifuge tubes with 0.1 M HCl. The absence of wall sorption in the experiment with low Ca(OH)2 concentration and pH = 10.3 suggests that this process is either dependent on pH or on the aqueous Ca concentration. The measured Np(V) concentration after centrifugation was similar to the one before centrifugation up to an added concentration of ~10-6 M (Fig. 24b). At starting Np(V) concentrations ≥ 10-6 M, a constant Np(V) concentration in solution of ~2·10-8 M was determined after centrifugation. This observation can be explained by the formation of Np(V) radiocolloids or a Np(V) solid phase, which was formed in this type of solution and settled during centrifugation. Note that the observed solution concentration is much higher than the Np(V) solubility limit deduced from speciation calculations which was based on the assumption that Ca0.5NpO2(OH)2·1.3H2O is the solubility limiting phase. Incomplete phase separation or formation of a poorly crystallized Ca-Np(V)-precipitate in these conditions are possible explanations of this observation.

The absence of Np(V) sorption on the walls of the centrifuge tubes in ACW (Fig. 24c) can be explained by the aqueous Np(V) speciation at this pH which is dominated by the anionic NpO2(OH)2

- species. Also in ACW, the measured Np(V) concentrations after centrifugation were similar to those before centrifugation up to an added concentration of ~10-6 M (Fig. 19c). At starting Np(V) concentrations ≥ 10-6 M, a constant Np(V) concentration in solution of approximately 10-6 M was determined after centrifugation. A closer look at the data reveals a slight difference between the aqueous Np(V) concentrations in 0.3 M NaOH and in ACW at


starting concentrations ≥ 10-6 M suggesting that the solubility was controlled by different solid phases. The measured aqueous Np(V) concentrations are again much higher than the Np(V) solubility limit calculated on the assumption of solubility control by Ca0.5NpO2(OH)2·1.3H2O.

The solubility tests carried out in this project confirm that the presence of Ca in the experimental systems considerably promotes the formation of Np(V) colloidal material which is likely to settle during centrifugation. In 0.3 M NaOH, the solubility tests indicate that Na-bearing Np(V) colloidal material was formed. This observation is consistent with the results of a study on the Np(V) solubility in dilute and concentrated alkaline NaCl solutions performed by Petrov et al. (2017) where the formation of Na-Np(V)-OH solid phases at Na concentrations > 0.1 M and above pH 13 was reported.

Fig. 24: Solubility tests of Np(V) in a) 10-4 M Ca(OH)2 at pH = 10.3, b) 10-2 M Ca(OH)2 at pH = 12.3, and c) 0.3 M NaOH and ACW. Np(V) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total Np(V) inventory. Equilibration time was 30 days. The line with slope +1 represents 100% recovery in solution.

10-8 10-7 10-6 10-510-8

10-7

10-6

10-5

10-4

Ca0.5NpO2(OH)2(s)

Before centrifugation After centrifugation

Mea

sure

d Np

(V) c

once

ntra

tion

(M)

Total Np(V) concentration (M)

10-4 M Ca(OH)2 (pH ≅ 10.3)


NpO2OH(am, fresh)

a

10-8 10-7 10-6 10-5 10-410-11

10-10

10-9

10-8

10-7

10-6

10-5

NpO2OH(am, fresh)

Ca0.5NpO2(OH)2(s)

Before centrifugationAfter centrifugation

Mea

sure

d Np

(V) c

once

ntra

tion

(M)


0.01 M Ca(OH)2 (pH=12.3)

b

10-8 10-7 10-6 10-5 10-410-11

10-10

10-9

10-8

10-7

10-6

10-5

10-4

0.3 M NaOH, before centrifugation 0.3 M NaOH, after centrifugation ACW, before centrifugation ACW, after centrifugation


Mea

sure

d Np

(V) c

once

ntra

tion

(M)



NpO2OH(am, fresh)

c


7.3 Sorption kinetics

The kinetics of the Np(V) sorption was determined on C–S–H phases with various C:S ratios in alkali-free solutions and on fresh HCP (Fig. 25). Sorption kinetic tests in ACW were not performed. The Rd,max value was estimated at 7.8·107 L kg-1.

The total Np(V) concentration used in these experiments was 5·10-6 M. The Np(V) concentrations in the supernatant solutions varied between 2·10-8 M < [Np] < 5·10-7 M. Comparison of the concentrations in the supernatant solution in the sorption experiment using a C–S–H phase having a C:S ratio of 0.65 with the results from the Np(V) solubility test conducted under similar chemical conditions (Fig. 24a) reveals that the solubility limits of the relevant Np(V) solids (NpO2OH(am, fresh) and Ca0.5NpO2(OH)2·1.3H2O(s)) were not exceeded in the sorption experiments. The aqueous Np(V) concentrations in the other sorption experiments were close to the experimental stability limits in similar chemical conditions (fresh HCP) or the thermodynamic solubility limits for Ca0.5NpO2(OH)2·1.3H2O(s) (C–S–H with C:S = 1.07 and 1.65). Thus, Np(V) precipitation could have occurred in these experiments. Note that, at the time when the experiments were performed, the existence of ternary Ca-Np(V)-OH solids had not yet been known. Experiments carried out at higher S:L ratios would have reduced the aqueous Np(V) concentrations below the solubility limits but also below the detection limit of LSA. Precipitation of Np(V) could be an artefact in these experiments. Further interpretation of the Np(V) sorption kinetics will therefore be based solely on the experimental data obtained with the C–S–H phase with a C:S ratio of 0.65. However, for the sake of comparison, all Np(V) sorption data are displayed.

Fig. 25: Np(V) sorption kinetics on C–S–H phases with three different C:S ratios in alkali-free conditions and on fresh HCP at pH 13.3. a) Aqueous Np(V) equilibrium concentrations versus time. b) Rd values versus time. Experimental conditions for the C–S–H experiments: S:L = 2·10-4 kg L-1, [Np]tot = 3.7·10-6 M, 2·10-8 M < [Np]eq < 10-7 M. Experimental conditions for the fresh HCP experiments in ACWHCP-I: S:L = 10-4 kg L-1, [Np]tot = 5·10-6 M, 2·10-7 M < [Np]eq < 5·10-7 M.

Sorption was found to be very fast. Very high Rd values of 105 L kg-1 were observed and attained within 1 day. The Rd values continued to increase slightly over a time period up to 10 days to reach a final steady state value of ~3·105 L kg-1. It is noteworthy that the Np(V) sorption appears to take significantly longer to attain steady state than the sorption of the trivalent and tetravalent actinides for which a steady state was reached usually within one day. However, the Np(V) sorption kinetics is still much faster than the recrystallization kinetics measured for the

0 10 20 30 40 5010-9

10-8

10-7

10-6

10-5

C:S = 0.65 C:S = 1.07 C:S = 1.65 Fresh HCP

Aque

ous N

p(V)

conc

entra

tions

(M

)

Time (days)

C-S-H phases in alkali-free conditions

a

0 10 20 30 40 50103

104

105

106

107

108

C:S = 0.65 C:S = 1.07 C:S = 1.65 Fresh HCP

R d(Np(

V))

(L k

g-1)

Time (days)


bRd,max


C–S–H phases indicating that recrystallization apparently has only a minor effect on the extent of Np(V) uptake. The slightly lower Rd value determined for Np(V) sorption onto fresh HCP (Fig. 25) can be explained on the assumption that C–S–H phases are the sole sorbing cement component and that C–S–H phases make up only ~50 wt% of fresh HCP (Lothenbach and Wieland, 2006).

7.4 Desorption tests

Desorption tests were performed to check for hysteresis of Np(V) sorption onto C–S–H phases: After sorbing Np(V) for a specific period of time, the solid and liquid phase were separated by centrifugation and the liquid phase was replaced by a new solution having a composition identical to the original liquid phase but without any Np(V). The Np(V) activity desorbing from the C–S–H phase and HCP into this new solution was determined as a function of time. The desorption experiment with the C–S–H phase with a C:S ratios of 0.65 in alkali-free conditions was set-up in such a way that the solubility limits of the relevant Np(V) solid phases (NpO2OH(am, fresh) and Ca0.5NpO2(OH)2·1.3H2O(s)) were not exceeded (2.0·10-7 M < [Np(V)]eq < 4.5·10-7 M). In the case of the desorption experiment with fresh HCP in ACW, however, the Np(V) supernatant concentration was expected to exceed the solubility limit of Ca0.5NpO2(OH)2·1.3H2O(s). The Rd,max value for these experiments was estimated at 1.33·109 L kg-1.

The experimental data show that the Rd values obtained from the desorption experiments after one day sorption and after 30 days sorption, rapidly return to the values observed during the original sorption tests (Fig. 26). This indicates that Np(V) sorption on this C–S–H phase is fully reversible and confirms the absence of an effect of a slow incorporation through recrystallization on the Rd values. The desorption data obtained from the experiment with HCP show a similar trend while it has to be noted that a portion of Np(V) could have been precipitated. This makes a thorough interpretation of the desorption data difficult.

Fig. 26: Np(V) desorption kinetics after 1 day sorption and 30 days or 60 days sorption. a) C–S–H phase with C:S ratio = 0.65 in alkali-free solution. b) fresh HCP in ACWHCP-I at pH = 13.3. Experimental conditions for the C–S–H system: S:L = 1.25·10-4 kg L-1, [Np]tot = 3.6·10-5 M, 10-7 M < [Np]eq < 4·10-7 M. Rd,max = 7·108 L kg-1. Experimental conditions for the HCP system: S:L = 1.25·10-4 kg L-1, [Np]tot = 3.6·10-5 M, 7·10-7 M < [Np]eq < 10-6 M.

0 10 20 30 40 50 60 70104

105

106

107

sorption Desorption after 1 day sorption Desorption after 30 days sorption

Rd(

Np(

V))

(L k

g-1)

Np(V) sorption onto C-S-H phases (C:S = 0.6). pH = 10.1

Time (days)

a

0 20 40 60 80 100 120104

105

106

107

Time (days)

Desorption after 1 day sorption Desorption after 60 days sorption Sorption

Rd(

Np(

V))

(L k

g-1)

b



Np(V) sorption isotherms on C–S–H phases with two different C:S ratios in alkali-free systems and in ACW were determined (Fig. 27). The slope of the isotherms in alkali-free solutions on a log-log scale is unity thus indicating linear sorption (Fig. 27a). Note that this linear behaviour of the Np(V) sorption onto the C–S–H phase with high C:S ratio extends far beyond the aqueous equilibrium Np(V) concentration (~2·10-8 M) at which the solubility tests showed precipitation of Ca-Np(V)-OH colloids (Fig. 24b). This discrepancy in the data has not yet been resolved. Two possible explanations are conceivable: 1) The presence of hitherto unknown ternary Np(V) hydroxy silicate complexes in the C–S–H solutions could increase the Ca0.5NpO2(OH)2·1.3H2O(s) solubility and thus allow Np(V) sorption experiments on these C–S–H phases to be conducted up to equilibrium concentrations far beyond the stabilities measured for the Ca-Np(V)-solid phase in alkaline solutions without Si. 2) Alternatively, the solubility limit could be shifted to higher Np(V) concentrations due to the formation of a more amorphous Ca-Np(V)-solid phase in these conditions.

Fig. 27: Np(V) sorption isotherms onto C–S–H phases with varying C:S ratios in alkali-free solution and in ACW at pH = 13.3. a, c) The amount of Np(V) sorbed versus the Np(V) aqueous equilibrium concentration. b, d) Rd(Np(V) versus the Np(V) aqueous equilibrium concentration. Experimental conditions: S:L = 5·10-3 kg L-1 (exp. 1) and 2·10-4 kg L-1 (exp. 2), equilibration time = 2 weeks. The red and black dashed vertical lines represent the experimental solubility limits determined previously (Fig. 24). The Rd,max values are shown as black dotted lines (Figs. b and d).

10-10 10-9 10-8 10-7 10-610-5

10-4

10-3

10-2

10-1

100

C:S=0.65: exp. 1 C:S=0.65: exp. 2 C:S=1.7: exp. 1 C:S=1.7: exp. 2


Np(V

) (so

rbed

) (m

ol k

g-1)

Np(V) equilibrium concentration (M)

a

10-10 10-9 10-8 10-7 10-6103

104

105

106

107


R d (N

p(V)

) (L

kg-1

)



Mean Rd = 3±2x105 L kg-1

Rd,max

b

10-11 10-10 10-9 10-8 10-7 10-610-5

10-4

10-3

10-2

10-1

100


ACW

Np(V

) (so

rbed

) (m

ol k

g-1)


Slope = 1

c

10-11 10-10 10-9 10-8 10-7 10-6104

105

106

107

108

C:S=0.65: exp. 1 C:S=0.65: exp. 2 C:S=1.25: exp. 1 C:S=1.25 exp. 2

R d (N

p(V)

) (L

kg-1

)


ACWRd,maxd


In ACW solutions, the Np(V) isotherm is characterized by a slope significantly less than unity on a log-log scale (Fig. 27c). This unusual sorption behaviour was observed in two independent series of sorption experiments. It may be an indication that more than one sorption site with different affinities for Np(V) could be involved. Such non-linear sorption behaviour, however, contradicts previous findings in this report suggesting that the sorption sites present on C–S–H phases are rather homogeneous w.r.t. their sorption properties and that the sorption capacity of the strongly sorbing sites is high. We presently don’t have an explanation for this unexpected behaviour.

7.6 Effect of pH and aqueous Ca concentration on the Np(V) sorption onto

TiO2

The Np(V) speciation calculations presented in Section 7.1 showed that the Np(V) speciation depends on the aqueous Ca concentration and on the pH in the range 10 < pH < 14. Sorption tests with Np(V) on TiO2 were carried out at different pH values to assess the potential effect of Ca onto Np(V) sorption. TiO2 is a solid phase known to be chemically stable in alkaline conditions and ambient temperature with low solubility and low recrystallization rate (Schmidt and Vogelsberger, 2009). Radionuclide incorporation into the structure of this mineral is unlikely, and therefore this solid is suited to study the surface complexation behaviour of radionuclides in alkaline conditions. Sorption experiments on TiO2 were carried out at a total Np concentration of ~1.3·10-10 M using a pure, short-lived 239Np tracer. Due to these extremely low 239Np concentrations, the formation of aqueous polynuclear species or solid phases can be excluded in these experiments.

Np(V) uptake by TiO2 was observed to decrease with increasing pH (Fig. 28a). The observed drop of the Rd value above pH = 12 corresponds very well with the increasing dominance of the NpO2(OH)2

- species in solution. This indicates that the affinity of this anionic species for the negatively charged TiO2 surface is very low, thus supporting the idea that anionic hydrolysis species stabilize the actinides in solution. On the basis of the concept of electrostatic inter ligand repulsion (Section 2.3), the limiting number of OH groups that fit in the equatorial plane of the NpO2

+ ion was determined to be nlimit = 2 (Fanghänel and Neck, 2002). This indicates that the anion NpO2(OH)2

- cannot bind to Ti-OH sorption sites due to coordinative saturation. To account for the effect of hydrolysis on sorption, a reduction factor, Fred, was defined based upon Eq. (3.14). This factor corresponds to the ratio of the distribution ratios in the absence (Rd

0) and in the presence (Rd) of the non-sorbing hydroxy species:

0 * 'd 2

redd

R KF 1R [H ]+

= = + , (7.1)

with *K2’ being the conditional stepwise stability constant for the reaction:

2 2 2 2NpO OH H O NpO (OH) H− ++ ⇔ + . (7.2)

The Rd values in the pH range 10 < pH < 14 can then be calculated as Rd = Rd0/Fred. It was

assumed that the Rd value determined at pH = 10 corresponds to Rd0. Log*K2

0 values at I = 0 are listed in Table 5. Ionic strength corrections for the calculation of the conditional stability constant, log*K2’, were carried out using the Davies equation (b = 0.3). The pH dependence of


the Rd values for Np(V) sorption on TiO2 can successfully be described with this model (Fig. 28a). This finding supports the idea that the NpO2(OH)2

- species is a limiting hydroxy complex which cannot form surface bonds with Ti-OH groups on TiO2.

Table 5: Overall and stepwise formation constants for relevant Np(V) hydrolysis complexes used in the model calculations (Log*Kn

0 = Log*βn0 - Log*βn-1).

Log*βn0 Log*Kn

0

NpVO2OH NpVO2(OH)2

- -11.3a)

-23.6a) -12.3

a) (Guillaumont et al., 2003)

The effect of the aqueous Ca concentration on the Np(V) uptake by TiO2 was studied in 0.3 M NaOH at pH = 13.3 ( Fig. 28b). Np(V) sorption was found to increase with increasing aqueous Ca concentration up to an equilibrium Ca concentration of ∼5·10−5 M. Ca is known to sorb on the surface of TiO2 (Tits et al., 2014b). The latter Ca concentration corresponds to maximum Ca site saturation on TiO2. Above this concentration, the sorbing Np(V) thus binds to a “Ca-OH” like surface rather than to a “Ti-OH” like surface. The effect of Ca sorption on Np(V) sorption can be explained either by neutralization of the negative surface charge, or by the formation of a very strong surface stabilized Ca-neptunate complex or Ca–neptunate ion pair. A thorough understanding of the Np(V) sorption processes on TiO2 under alkaline conditions in the presence of Ca would require a more in-depth wet chemistry and spectroscopic investigation. On the basis of the current results we can conclude that both the pH and the Ca concentration have a significant influence of the Np(V) surface complexation on TiO2.

Fig. 28: Np(V) sorption onto TiO2. Effect of pHc (a) and of the aqueous Ca2+ concentration in ACW (b). The solid line in (a) represents the model calculation using the equation Rd = Rd

0/Fred. The solid line in (b) is added to guide the eye.

7.7 Effect of the C:S ratio of C–S–H phases

The C:S ratio of C–S–H phases is known to correlate with both the pH and Ca concentration in solution (see Section 4.2). No effect of the C:S ratio on the Np(V) sorption onto C–S–H phases was observed in the sorption kinetic experiments and the sorption isotherms described earlier

10 11 12 13 14 15103

104

105

106

107

R d (N

p(V)

) (L

kg-1

)

pHc

a

R0d

10-6 10-5 10-4 10-30104

105

106

107

R d (N

p(V)

(L

kg-1

)

Ca2+ concentration (M)

ACW, pH=13.3b


contrary to our expectations. To confirm this observation, a detailed study of the effect of the C:S ratio on the Np(V) sorption was carried out (Fig. 29). Note that the Rd,max value was estimated to be 7·107 L kg-1 in this experiment. Note further that the short half-life of the 239Np tracer (2.36 days) that was used for these experiments prompted us to limit the equilibration time to three days. Thus, it is acknowledged that in some cases steady state may not have been attained within this short equilibration time. Nevertheless, we believe that this uncertainty might be smaller than the overall experimental uncertainty on the measurements. The Rd values for the C–S–H phases with increasing C:S ratio in alkali-free systems and in ACW (pH 13.3) were found to be constant (Fig. 29a). Reversible adsorption onto C–S–H phases, similar to that on TiO2, would imply that, with increasing C:S ratio, the Rd values are affected simultaneously by two counteracting parameters: 1) a decrease in Rd value with increasing pH due to increasing concentrations of the non-sorbing anionic NpO2(OH)2

− complex in solution, and 2) an increase in Rd value with increasing aqueous Ca concentration. Both effects may, at least partially, compensate each other, which complicates a direct comparison of the sorption behaviour on TiO2 and C–S–H phases. The single effect of the aqueous Ca concentration at constant pH on Np(V) uptake by C–S–H phases is apparent from the sorption tests in the presence of alkalis (ACW, pH = 13.3) (Fig. 29c). In contrast to Np(V) sorption onto TiO2, no effect of an increase in the aqueous Ca concentration from 10−5 M to 10−3 M on the Np(V) sorption onto C–S–H phases was observed. This observation suggests that other processes different from surface complexation are controlling the Np(V) sorption onto C–S–H phases and that the Np(V) sorption mechanisms onto C–S–H phases are still not sufficiently understood.

7.8 Discussion

7.8.1 Sorption data The sorption behaviour of Np(V) on C–S–H phases and HCP can be summarized as follows:

1) Very strong sorption illustrated by Rd values in the range of 105 L kg-1 < Rd < 107 L kg-1;

2) Relatively fast sorption kinetics onto C–S–H phases and HCP (steady state in < 7 days), while slower than the sorption kinetics of other actinides;

3) Reversible sorption process; 4) Linear sorption under alkali-free conditions but non-linear in ACW at pH = 13.3; 5) Decreasing Np(V) sorption with increasing pH on TiO2 due to the formation of the non-

sorbing NpO2(OH)2- species;

6) Increasing sorption on TiO2 with increasing Ca2+ concentrations in solution; 7) No effect of the C–S–H composition (C:S ratio) (i.e. aqueous Ca2+ concentration and

pH) on Np(V) sorption.

There are not many experimental studies published in the open literature on the sorption of Np(V) onto cementitious materials. In their literature review of radionuclide sorption onto cement, Ochs et al. (2016) listed experimental data from studies on different cements published by Höglund et al. (1985) and Allard et al. (1985). The sorption data listed in the review are very high (up to 106 L kg-1) and tend to decrease with increasing pH from 5·103 –106 L kg-1 at pH = 12, to ~103 – 3·104 L kg-1 at pH 13.3.


Fig. 29: Sorption of Np(V) on C–S–H phases in alkali-free conditions (10 < pH < 12.5) and in ACW at a constant pH of 13.3. a) Effect of the C:S ratio in alkali-free conditions. b) Rd values from a) plotted versus pHc. c) Rd values from a) measured in ACW at fixed pH = 13.3 plotted versus the free Ca2+ concentration

The influence of pH reported by Ochs et al. (2016) is in contrast to the pH independent sorption observed in the present study. Note, however, that the sorption experiments on TiO2 have shown that the aqueous Ca concentration also plays a role on the Np(V) sorption and this parameter was not discussed in Ochs et al. (2016). In addition, presently unknown ternary Np(V) hydroxyl silicate complexes in solution may also play a role in the Np(V) sorption behaviour.

The strong sorption of Np(V) on TiO2 at pH = 10 and on C–S–H phases, observed in the present experiments and partially in the literature, was not anticipated as the low effective charge (2.2, NpO2

+) rather suggests weak to moderate sorption (see Section 2.3). Ochs et al. (2016) and Sylwester et al. (2000) argued that these high Rd values may originate from partial reduction of Np(V) after sorption onto the cementitious materials. This argument, however, ignores the results from the XAS studies by Gaona et al. (2013) which did not indicate a reduction process in C–S–H and HCP systems.

Np(V) sorption data reported for various types of solids at pH = 10 are controversial. Measured Rd values were found to vary from 1.0 - 400 L kg-1 on quartz (Bertetti et al., 1998; Kohler et al., 1999; Pathak and Choppin, 2007), 4·102 - 3·104 L kg-1 on clay minerals (Gorgeon, 1994; Bertetti et al., 1998; Turner et al., 1998; Amayri et al., 2011) and 8·103 - 105 L kg-1 on Fe-oxides (Kohler et al., 1999; Khasanova et al., 2007). This large variability in the Rd (or Kd, respectively) values could be explained partially by sorbent properties such as the effect of

0.6 0.8 1.0 1.2 1.4 1.6 1.8102

103

104

105

106

107

Alkali-free ACW

R d (N

p(V)

) (L

kg-1

)

C:S (mol/mol)

a

10 11 12 13 14102

103

104

105

106

107

Alkali-free ACW

R d (N

p(V)

) (L

kg-1

)

pHc

b

R0d=5·105 L Kg-1

10-5 10-4 10-3104

105

106

107

R d (L

kg-1

)

Ca2+ equilibrium concentration (M)

c ACW


surface loading (non-linear sorption) and sorption capacity of the different solids (Bertetti et al., 1998). Furthermore, the type of Np(V) surface complexes formed (monodentate versus multidentate) could influence measured Rd values as well. The high Rd values determined for Np(V) sorption on C–S–H phases are supported by similarily high Rd values determined for Np(V) on oxides at pH = 10 (Kohler et al., 1999; Khasanova et al., 2007).

More qualitative support for the high Np(V) sorption values can be obtained by comparing the Rd values for the sorption of actinides in different redox states onto C–S–H phases with their respective stepwise hydrolysis constants (stability constant for the addition of one OH group to the relevant actinide species) at pH = 10. The assumption underlying this comparison is that the formation of an actinide surface complex (Eq. 7.3) and the stepwise hydrolysis of an actinide (Eq. 7.7) are chemically similar processes and that “Rd” is a measure for the strength of the actinide surface complex (Eq. 7.6) and thus that “Rd” and *Kn+1

0 are correlated.

The surface complexation of an actinide can be written as follows:

z n z (n 1)n nSOH An(OH) SO An(OH) H− − + +≡ + ⇔≡ − + (7.3)

with a surface complexation constant as follwos:

z (n 1)

An nsurf z n

n

[ SO An(OH) ][H ]K[ SOH][An(OH) ]

− + +

−≡ −

=≡

(7.4)

Based upon the definition of the Rd value in Eq. (3.4), in this specific case, Rd can be written as:

z (n 1)n

d z nn

[ SO An(OH) ] VRm[An(OH) ]

− +

−≡ −

= ⋅ (7.5)

Combining Eqs. (7.4) and (7.5) gives:

And surf

[ SOH] VR Km[H ]+

≡= ⋅ ⋅ (7.6)

Assuming a constant pH ([H+] is constant) and trace An concentrations ([ SOH]≡ is constant), Rd and KAn

surf are correlated and Rd is thus a measure of the strength of the surface complex.

The stepwise hydrolysis reaction of an actinide can be written in a similar form as the surface complexation reaction in Eq. (7.3):

z (n 1)z nn 2 n 1An(OH) H O An(OH) H− +− +

++ ⇔ + , (7.7)

with the corresponding conditional stepwise surface complexation constant, *Kn+1’:

' z (n 1) z nn 1 n 1 n*K An(OH) H An(OH)− + + −

+ + = . (7.8)

' ' 'n 1 n 1 nlog*K log* log*+ += β − β (7.9)


The conditional stability constant for the relevant stepwise hydrolysis reaction of Np(V) is very similar to the relevant stepwise hydrolysis constants of the other actinides listed in Table 6 (except for Am(III)). Accepting the similarity between surface complexation reactions and stepwise hydrolysis reactions described above, this provides support for the observation that the Rd values for the sorption of these actinides are also similar at pH = 10.

Table 6: Stepwise hydrolysis constants for different actinides and lanthanides at pH = 10. Comparison with the log Rd values of the respective actinides and lanthanides.

Actinide

/lanthanide

Stepwise hydrolysis reaction * 'n 1log K + log Rd

(at pH = 10.0)

L kg-1

EuIII(OH)4-

AmIII(OH)4-

ThIV(OH)5-

NpVO2(OH)2-

UVIO2(OH)42-

NpVIO2(OH)42-

3 2 4Eu(OH) (aq) H O Eu(OH) H− ++ ⇔ +

3 2 4Am(OH) (aq) H O Am(OH) H− ++ ⇔ +

4 2 5Th(OH) (aq) H O Th(OH) H− ++ ⇔ +

2 2 22NpO (OH) H O NpO (OH) H− ++ ⇔ +

23 2 22 4UO (OH) H O UO (OH) H− − ++ ⇔ +

23 2 22 4NpO (OH) H O NpO (OH) H− − ++ ⇔ +

-12.52

-14.52

?

-12.32

-11.22

-10.82

~5.5

~5.5

~5.9

~5.7

~6.01

~6.01 1Sorption data discussed in the following section 2Calculated from the stability constants listed in the Tables 3, 4, 5 and 7

7.8.2 Uptake mechanisms The XAS investigations published by Gaona et al. (2013b) indicate that adsorbed Np(V) could be accommodated in the interlayer space of the C–S–H phases. Indeed, upon sorption, longer Np(V)-Oax distances and shorter Np(V)-Oeq distances were observed compared to the NpVO2

+ aquo ion. This implies the removal of most of the H2O molecules from the first coordination sphere of the sorbed Np(V) indicating that a significant portion of the sorbed Np(V) is accommodated in the C–S–H interlayer. The absence of any effect of the aqueous Ca concentration on the Np(V) uptake by C–S–H phases further shows that the uptake mechanism into the C–S–H interlayer is not simply due to a substitution of Ca(II) by Np(V) in the interlayer (Fig. 29c). On the contrary, the presence of Ca seems to stabilise sorbed Np(V) species as shown by the sorption experiments onto TiO2 (Fig. 22b). The XAS investigations of Gaona and co-workers did not confirm the presence of Ca coordinated to the sorbing Np(V). The backscattering contributions from neighbouring atoms beyond the axial and equatorial oxygen shells were found to be weak and attributed to neighbouring Si atoms rather than Ca atoms. Thus, the exact Np(V) uptake mechanism cannot be unraveled from the currently available wet chemistry and spectroscopic information.

The desorption studies with Np(V) on C–S–H and HCP indicate that Np(V) uptake by C–S–H phases is a reversible process. Difference in the Np(V) sorption kinetics and C–S–H recrystallization kinetics suggest that Np(V) uptake into the interlayer occurs via diffusion into the C–S–H interlayers rather than recrystallization.


8 Sorption of hexavalent actinides

8.1 Speciation

Actinides having stable hexavalent oxidation states are mainly U, Np and Pu, forming linear dioxo moieties (O=An=O)2+. The present work focusses on the sorption of U(VI) and Np(VI). Speciation calculations were performed with the code “Medusa” (Puigdomenech, 1983). The thermodynamic complexation constants were taken from Thoenen et al. (2014) for the silicate complex and the from the more recent publications of Gaona et al (2013a) and Altmaier et al (2017) for the U(VI) and Np(VI) hydroxy complexes and Na2U2O7·H2O(cr) solubility product. The thermodynamic data used are listed in Table 7. Only the An(VI) hydroxy complexes were included in the calculations. Thermodynamic data for ternary Ca-hydroxy-uranyl complexes are not yet available. An(VI)-silicate complexation is only known at pH < 5. Thermodynamic complexation constants for all other reactions needed in the calculations were taken from Hummel et al. (2002). All calculations were performed fixing I to 0.3 M in the code Medusa. This value reflects the ionic strength in ACW. Note that this implies that the speciation calculated at pH > 13.3 is only approximately correct.

Table 7: Summary of log10*β° and log10*K°s,0 used in the An(VI) speciation calculations (An(VI) = U(VI) or Np(VI)). Only species formed between 10 < pH < 14 are considered (Gaona et al. (2013a), Thoenen et al. (2014), Altmaier et al. (2017)).

Species Reaction log10*β°

U(VI)

log10*β°

Np(VI)

AnO2(OH)3-

AnO2(OH)42-

AnO2SiO3(OH)3+

22 2 32AnO 3H O AnO (OH) 3H+ − ++ ⇔ +

2 22 22 4AnO 4H O AnO (OH) 4H+ − ++ ⇔ +

23 22 3AnO SiO(OH) AnO SiO(OH)+ − ++ ⇔

-20.7

-31.9

+7.8

-21.2

-32.0

Solid Reaction log10*K°s,0 log10

*K°s,0

AnO2(OH)2·H2O(cr)

Na2An2O7(cr)

CaAnO4(s)

22 2 2 22AnO 3H O AnO (OH) ·H O(cr) 2H+ ++ ⇔ +

22 2 2 722Na 2AnO 3H O Na An O (cr) 6H+ + ++ + ⇔ +

2 22 42Ca AnO 2H O CaAnO (s) 4H+ + ++ + ⇔ +

-5.3

-22.6

-23.1

-5.47

-25.1

The speciation of U(VI) and Np(VI) in alkaline conditions between pH 10 and 13.3 in the absence of Ca is characterized by the formation of the two dominant hydrolysis products: AnO2(OH)3

- and AnO2(OH)42- (see Fig. 30a for the case of U(VI)). The speciation of U(VI) and

Np(VI) are very similar. The solubility controlling An(VI) phases were considered to be Na2U2O7·H2O(cr) (Altmaier et al., 2017) and Na2Np2O7(cr) (Gaona et al., 2013a) in Na rich solutions. The solubilities of the An(VI) containing solids increase with increasing pH due to the increasing stabilization of the An2(OH)4

2- complex in solution.


Fig. 30: U(VI) speciation calculations in the pH range 10 < pH < 14 in the absence of Ca (a) and in the presence of high Ca concentrations (b) representing the maximal Ca concentration in C–S–H solutions in the absence of alkalis.

The aqueous speciation of the hexavalent actinides under alkaline conditions in the presence of Ca is still not well explored. It can, however, be assumed that the hexavalent actinides form ternary Ca-An(VI)-OH complexes comparable to the tri-, tetra-and pentavalent actinides. Ca-uranates are potential solubility-limiting U(VI) solid phases under alkaline conditions in the presence of Ca. Becquerelite is the only well-characterized Ca-uranate for which sufficiently detailed solubility measurements are available (Guillaumont et al., 2003). In case of the other Ca-uranates taken into account by Guillaumont et al. (2003), the solubility constants were mainly obtained from calorimetric experiments. Hummel and co-workers (Hummel et al., 2002) reject these data arguing that solid phases identified and characterized in calorimetric experiments at high temperature are often not representative for aqueous systems. CaUO4(s) is the only Ca-uranate given by Guillaumont et al. (2003) for which solubility data were deduced from wet chemistry experiments. This phase was identified and characterized in a study on the uranium solid speciation in cementitious materials reported by Moroni and Glasser (1995). Using the experimental data reported by Moroni and Glasser (1995), Guillaumont et al. (2003) estimated the solubility product for the following reaction:

2 22 42Ca UO 2H O CaUO (s) 4H+ + ++ + ⇔ + 0

10 s,0log K = 23.1 (8.1)

Note that this Ca-uranate phase was not selected in the thermodynamic database of Guillaumont et al. (2003) because the experimental system was not completely characterized in the original work and the 0

10 s,0log K derived can only be seen as a rough estimate of the solubility product of CaUO4(s). The solubilities of Ca-neptunate phases are not known, but a solubility behaviour similar to that of CaUO4(s) is anticipated. It is worth mentioning that, very recently, Fellhauer et al. (Fellhauer et al., 2018) published new data on the Np(VI) solubility and the formation of calcium neptunates in alkaline CaCl2 solutions. Unfortunately, the publication of these data came too late to be considered in the solubility calculations carried out in the present report.

Fig. 30b shows the results of U(VI) solubility and speciation calculations under alkaline conditions in the presence of 0.02 M Ca, i.e. the maximal Ca concentration expected in solutions in equilibrium with C–S–H phases in the absence of alkalis. Note that the total U(VI) concentration in this calculation was fixed at values below the total Ca concentrations (realistic assumption) but still high enough to allow precipitation of solid U(VI) phases. The U(VI) solubility decreases with increasing pH due to the increasing stabilization of the CaUO4(s) solid compared to the UO2(OH)3

- complex. U(VI) solubility reaches a minimum in the pH region

10 11 12 13 14-10

-9

-8

-7

-6

-5

-4

-3a

UO2(OH)2-4

UO2(OH)-3

Log[

U(V

I)] (M

)

pH

Na2U2O7·H2O(cr)

[UO2+2 ]tot=10-5 M

[Ca]tot= 0 M[Na]tot = 0.3 M

10 11 12 13 14-8

-7

-6

-5

-4

-3b

UO2(OH)2-4

UO2(OH)-3

Log

[U(V

I)] (

M)

pH

[UO2+2 ]tot=10-5M

[Ca]tot = 0.02 M[Na]tot = 0 M

CaUO4(s)

Ca(OH)2

precipitation


where the UO2(OH)42- species dominates the aqueous U(VI) speciation. Note that the speciation

calculations in the presence of Ca have to be taken with great caution as aqueous ternary Ca-An(VI)-OH complexes are not included in these calculations. It may be assumed that the Np(VI) speciation is quite similar to the U(VI) speciation.

The An(VI)-silicate complexes prevailing in the low pH region (AnO2SiO(OH)3+) do not play a

significant role in the An(VI) speciation in alkaline conditions. Stability constants for the ternary An(VI)–hydroxy–silicate complexes that could exist in the alkaline pH region are not yet available.


The previous discussion of the An(VI) speciation under cement pore water conditions in the presence of Ca show that key thermodynamic information, e.g. on the existence and stability of ternary Ca-An(VI)-OH complexes, on the solubility of alkali- and alkaline-earth uranates/neptunates and also on the existence and stability of ternary An(VI)-hydroxy-silicate complexes is still lacking. Therefore, experimental checks of the stability of U(VI) in alkaline solutions with Ca concentrations similar to the concentrations in C–S–H solutions and cement pore water solutions were carried out prior to performing sorption experiments. The effect of silicate concentrations on the stability of U(VI) in alkaline solutions was not recognised at the time these experiments were planned and was therefore not tested. Solubility tests with Np(VI) were not performed because it was assumed that its speciation is largely identical to that of U(VI). The solubility tests in alkali-free solutions and in ACW at pH = 13.3 show the ratio of the radionuclide concentration determined in solution to the total added radionuclide concentration (Figs. 31a, 31c, 31e, 32a) and the measured U(VI) solution concentration as a function of the added U(VI) concentration in solution (Figs. 31b, 31d, 31f, 32a). The data refer to measured concentrations before centrifugation (1 h at 95,000g) and after centrifugation. Note that particles with diameter larger than ~3 nm are expected to settle during centrifugation. Thus, the concentration determined after centrifugation represents the concentration of dissolved radionuclides plus the concentration of radionuclides associated with colloidal material smaller than ~3 nm.

The results from the solubility test carried out in alkali-free solutions in the presence of 5·10-5 M Ca (Figs. 31a and b) show no U(VI) reduction by settling of U(VI) colloids after centrifugation and an effect of the equilibration time was not observed. The U(VI) solutions were stable in the concentration range covered in the experiments, i.e. up to a total U(VI) concentration of 10-4 M. Figs. 31c and d further reveal that in an U(VI) solution with a Ca concentration of 3·10-3 M no reduction of the U(VI) concentration was observed before and after centrifugation up to a total U(VI) concentration of 10-5 M. Above this concentration threshold, a small reduction was observed before centrifugation. This reduction was more pronounced after centrifugation (≥ 60%), suggesting the presence of colloidal U(VI) solid phases. Fig. 31d further indicates that U(VI) solutions were stable in the presence of 3·10-3 M Ca up to a total U(VI) concentration of approximately 10-5 M. The effect of the equilibration time on U(VI) solution stabilities was negligible. In a 0.015 M Ca(OH)2 solution, the U(VI) concentrations determined in solution before and after centrifugation were between 10% and 20% lower than the initial U(VI) concentration below a total U(VI) concentration of 2·10-6 M (Figs. 31e and f). This indicates that a portion of the U(VI) was adsorbed onto the walls of the centrifuge tubes. The missing U(VI) could be recovered by washing the centrifuge tubes with 0.1 M HCl. Significant activity reduction was observed above a concentration of 2·10-6 M both in the samples measured before and after centrifugation. This reduction in the U(VI) concentration was attributed to the formation of colloidal U(VI) solid phases. The observed reduction before centrifugation were


caused by larger U(VI) colloids, which already settled before centrifugation. The activity reduction observed after centrifugation is attributed to the presence of smaller colloids suspended in solution, which settled during centrifugation. Based on these tests it was concluded that U(VI) solutions containing 0.015 M Ca were stable up to an U(VI) concentration of about 2·10-6 M. In these experiments as well, the effect of the equilibration time (1 day and 7 days) on U(VI) solution stabilities was negligible.

Fig. 31: Solubility tests with U(VI) in alkali-free solutions. Figs. a, c, e: Percentage of the U(VI) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Figs. b, d, f: U(VI) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Equilibration time was 1 day and 7 days in alkali-free systems. The curves in Figs. a, c, e only serve to guide the eye. The lines with slope +1 in Figs. b, d, f represent 100% recovery in solution.

10-8 10-7 10-6 10-5 10-4 10-30

20

40

60

80

100

120[Ca] = 5·10-5 MpH = 10.1


1 day, before centr. 1 day, after centr. 7 days, before centr. 7 days, after centr.%

U(V

I) m

easu

red

in so

lutio

n

Total U(VI) concentration (M)

a

10-8 10-7 10-6 10-5 10-4 10-310-8

10-7

10-6

10-5

10-4

10-3

1 day, before centr. 1 day, after centr. 7 days, before centr. 7 days, after centr.

U(VI

) con

cent

ratio

n in

solu

tion

(MTotal U(VI) concentration (M)

[Ca] = 5·10-5 MpH = 10.1

b Alkali-free conditions

10-8 10-7 10-6 10-5 10-4 10-30

20

40

60

80

100

120 [Ca]=3·10-3 MpH = 12.0

1 day, before centr. 1 day, after centr. 7 days, before centr. 7 days, after centr.%

U(V

I) m

easu

red

in so

lutio

n


c Alkali-free conditions

10-8 10-7 10-6 10-5 10-4 10-310-8

10-7

10-6

10-5

10-4

10-3Alkali-free conditions


U(VI

) con

cent

ratio

n in

solu

tion

(M)


[Ca]=3·10-3 MpH = 12.0

d

10-8 10-7 10-6 10-5 10-4 10-30

20

40

60

80

100

120[Ca] = 1.5·10-2 MpH = 12.4


% U

(VI)

mea

sure

d in

solu

tion


e Alkali-free conditions

10-8 10-7 10-6 10-5 10-4 10-310-8

10-7

10-6

10-5

10-4

10-3Alkali-free conditions

U(VI

) con

cent

ratio

n in

solu

tion

(M)


[Ca] = 1.5·10-2 MpH = 12.4

Stability limit

CaUO4(s) 1 day, before centr. 1 day, after centr. 7 days, before centr. 7 days, after centr.

f


The above tests show that the experimental solubility limit of U(VI) solutions in the pH range between 12 - 12.5 increases with decreasing Ca concentrations suggesting that Ca-uranate – type phases control the U(VI) solubility.

In the case of the U(VI)-ACW system no reduction in the U(VI) concentration due to wall sorption was observed. Furthermore, the difference in the concentrations after equilibration for 1 day and 32 days was negligible. No reduction in the solution concentrations were observed before and after centrifugation up to a starting concentration of about 2·10-6 M (Fig. 32a). The data fall on the line with slope = 1 (Fig. 32b). Above this concentration threshold, however, the portion of added U(VI) that remained in solution after equilibration was found to decrease presumably due to the formation of a U(VI) solid phase. This decrease was more pronounced after centrifugation (Fig. 32a). The U(VI) concentration in solution was found to be constant at about 10-5 M. Note that this concentration is identical to the solubility limit of Na2U2O7·H2O(cr) found by Altmaier et al. (2017) (Fig. 32b). The speciation calculations indicate that the latter phase could limit U(VI) solubility in ACW. Based on the above findings it was considered that the initial U(VI) concentration in ACW in sorption experiments should not exceed 7·10-6 M to minimize the risk of artefacts due to Na2U2O7(s) precipitation.

Fig. 32: Solubility tests with U(VI) in ACW. a) Percentage of the U(VI) inventory measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. b) U(VI) concentration measured in solution before (open symbols) and after (closed symbols) centrifugation as function of the total U(VI) inventory. Equilibration time was 1 day and 32 days. The curves in Fig. a only serve to guide the eye. The lines with slope +1 in Fig. b represent 100% recovery in solution.

8.3 Sorption kinetics

U(VI) sorption kinetic experiments with C–S–H phases in ACW and in alkali-free solutions (pH 11.4 - 12.5) show that the Rd values determined for the different C–S–H phases ranged in value between ~102 L kg-1 in ACW and 106 L kg-1 in alkali-free solution. Note that the Rd,max was estimated at 1.5·106 L kg-1 for all the experiments. U(VI) uptake was stronger in alkali-free solution, in the pH range between 11.4 and 12.5 than in ACW (pH = 13.3). Time dependence of the Rd values was similar in ACW and in alkali-free solutions. Sorption was found to proceed in two steps: after a very fast initial uptake step taking ~1 day, Rd values continued to increase slightly to attain steady state after ~20 days equilibration in most of the experiments. Similar uptake kinetics were also observed in previous studies by Pointeau et al. (2004b) and suggest that the initial uptake mechanism involves a surface sorption process followed by slow

10-8 10-7 10-6 10-5 10-4 10-30

20

40

60

80

100

ACW


% U

(VI)

mea

sure

d in

solu

tion


a [Ca]=1.6·10-3 M

10-8 10-7 10-6 10-5 10-4 10-310-9

10-8

10-7

10-6

10-5

10-4

10-3[Ca]=1.6·10-3 M

32 days, before centr. 32 days, after centr. 1 day, before centr. 1 day, after centr

Mea

sure

d U(

VI) c

once

ntra

tion

(M)


Na2U2O7·H2O(cr) solubility limit

b

Slope = +1

ACW


incorporation into the structure. Note that this sorption kinetic behaviour is very similar to that of the pentavalent actinides but different from the sorption kinetics observed before for the trivalent and tetravalent actinides/lanthanides. Sorption reactions of the latter actinides reached steady state within one day without showing a second slow sorption step. In ACW, Rd values were found to increase with increasing C:S ratio. However, this effect of the C:S ratio could not be confirmed in alkali-free systems.

The uptake behaviour observed in the co-precipitation experiments was similar to that in the sorption experiments (Fig. 33b). Note, however, that the Rd values obtained from the co-precipitation experiments in ACW were significantly higher than those determined in the sorption experiments. The difference in Rd values was about one order of a magnitude for the same C:S ratios. By contrast, no difference in the Rd values deduced from the sorption and co-precipitation experiments was observed in alkali-free solution within the uncertainty on the data. At present, this observation cannot be explained conclusively.

Fig. 33: Sorption kinetics (a) and co-precipitation kinetics (b) of U(VI) by C–S–H phases with different C:S ratios in alkali-free solutions (pH between 10.0 and 12.5) and in ACW (pH 13.3). Experimental conditions: S:L ratio = 5·10-3 kg L-1, [U]tot = 5.4·10-7 M, Rd,max = 1.5·106 L kg-1.

The relative uncertainty on the U(VI) Rd values in the sorption and co-precipitation experiments was estimated to be 50%. These high uncertainties are mainly related to the very strong uptake, which results in ratios of the sorbed activity to the activity in the aqueous phase larger than

0 20 40 60 80 100 120101

102

103

104

105

106 C:S=0.75; Alkali-free C:S=1.07; Alkali-free C:S=1.65; Alkali-free C:S = 0.75; ACW C:S = 1.07;ACW C:S=1.25; ACW

R d(U

(VI))

(L

kg-1

)


a

0 20 40 60 80 100 120103

104

105

106

107 C:S=0.7; Alkali-free C:S=1.07; Alkali-free C:S=1.25; Alkali-free C:S=1.65; Alkali-free C:S = 0.75; ACW C:S=1.07; ACW

R d(U

(VI))

(L

kg-1

)


b


1000. Incomplete phase separation then may easily lead to significant overestimation of the 233U activities in the aqueous phase.

U(VI) sorption kinetic data on fresh HCP in ACW are comparable with those on the C–S–H phase with a C:S ratio of 1.25 in ACW. Both sets of data are consistent with a view to the uncertainties on the measurements (see Fig. 34).

Fig. 34: Sorption kinetics of U(VI) onto fresh HCP in ACWHCP-I (pH 13.3). Experimental conditions: S:L = 10-3 kg L-1, [U]tot = 5.3·10-7 M.

8.4 Effect of the S:L ratio

Sorption studies with Eu(III) showed that the measured Rd values may depend on the S:L ratio of the experimental system due to incomplete phase separation. In the case of U(VI) a potential effect of the S:L ratio was checked only on fresh HCP (Fig. 35). Two identical experiments were carried out. The experimental data clearly show that the S:L ratio of the HCP suspensions does not have a significant effect on the sorption of U(VI).

Fig. 35: U(VI) sorption as a function of the S:L ratio in ACWHCP-I at pH = 13.3 on fresh HCP. Experimental conditions: [U(VI)]tot = 5.4·10-7 M; equilibration time = 30 days.

0 10 20 30 40 50 60 70 80102

103

104

105

No pre-equilibration After pre-equilibration; exp. 1 After pre-equilibration; exp. 2

R d (U

(VI))

(L k

g-1)

Time (days)

Fresh HCP

10-5 10-4 10-3 10-2 10-1102

103

104

Exp. 1 Exp. 2

Rd (

U(V

I)) (

L kg

-1)

S:L (kg L-1)



U(VI) isotherms on C–S–H phases with two different C:S ratios in alkali-free systems and in ACW are shown together with the corresponding plots of the Rd behaviour (Fig. 36). Rd,max values (thin dotted lines in Fig. 36b) are given for the isotherm measured onto a C–S–H phase with C:S ratio of 0.75 because this was the only dataset for which the measured Rd values could be close to the Rd,max.

In the alkali-free systems the experimental U(VI) stability limits in solution were never exceeded. In ACW, the highest experimental points of the isotherms coincided with the measured U(VI) stability limits in solution.

The sorption isotherms were generally highly non-linear both under alkali-free conditions ( 10 < pH < 12.5) and in ACW (pH 13.3) indicating that sites with different affinities for U(VI) are present in the C–S–H phases. Furthermore, U(VI) sorption clearly depends on the C–S–H composition and on pH. In ACW (constant pH), the U(VI) sorption increased with increasing C:S ratio and thus with increasing aqueous Ca concentration. The presence of Ca thus appears to facilitate U(VI) uptake. Comparison of the sorption data for the same C–S–H phase in alkali-free solution (10.1 < pH < 12.5) and in presence of ACW (pH 13.3) shows that U(VI) sorption decreased with increasing pH. This pH dependence is in agreement with observations made by Pointeau et al. (2004b) on U(VI) uptake by C–S–H phases and degraded cement pastes and is explained by a change of the aqueous U(VI) speciation. Indeed, with pH increasing from 10 to 13.3, the U(VI) speciation changes from an UO2(OH)3

- dominated to an UO2(OH)42- dominated

speciation. Thus, with increasing pH, equilibria are progressively shifted towards the aqueous hydroxyl species at the expense of the sorbed species. The Ca, Si, Na and K concentrations were constant over the entire range of U(VI) loadings for all the sorption isotherms (data not shown), indicating that the uptake of U(VI) did not involve a measurable release of other cations or gave rise to the transformation of the C–S–H phases.

Fig. 36: U(VI) sorption isotherms on C–S–H phases with different Ca:Si ratios in alkali-free conditions and in ACW at pH = 13.3. a) The amount of U(VI) sorbed versus the U(VI) aqueous equilibrium concentration. b) Rd versus the U(VI) aqueous equilibrium concentration. Vertical grey bars represent the solubility of Na–uranate in the ACW. Dotted lines represent the Rd,max values. Dashed and solid lines are added to guide the eye.

A sorption U(VI) isotherm measured on fresh HCP in ACW at pH = 13.3 corresponds to the sorption isotherm measured on a C–S–H phase with a C:S ratio of 1.25 in ACW (Fig. 37).

10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-410-6

10-5

10-4

10-3

10-2

10-1

C:S = 0.75; alkali-free, pH=10.1 C:S = 1.07; alkali-free, pH=12.1 C:S = 1.65; alkali-free, pH=12.5 C:S = 0.74 ACW, pH=13.3 C:S = 1.07 ACW, pH=13.3 C:S = 1.25 ACW, pH=13.3

U(VI

) sorb

ed (m

ol k

g-1)

U(VI) equilibrium concentration (M)

a

Slope = +1

U(VI) stability limit in ACW

10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4101

102

103

104

105

106

107

C:S = 0.75; alkali-free, pH=10.1 C:S = 1.07; alkali-free, pH=12.1 C:S = 1.65; alkali-free, pH=12.5 C:S = 0.74 ACW, pH=13.3 C:S = 1.07 ACW, pH=13.3 C:S = 1.25 ACW, pH=13.3

R d (U(V

I)) (

L kg

-1)


b

U(VI) stability limit in ACW


U(VI) sorption onto HCP was found to be slightly non-linear but the non-linearity was less than in the experiments with C–S–H phases. U(VI) precipitation effects were observed at an equilibrium U(VI) concentration of ~10-5 M in agreement with the experimental solubility tests described above (Fig. 31h).

Fig. 37: U(VI) sorption isotherms on fresh HCP in ACWHCP-I (pH 13.3). a) The amount of U(VI) sorbed versus the aqueous equilibrium concentration. b) Rd versus the aqueous equilibrium concentration. Thermodynamic U(VI) solubility taken from Fig. 30b. Experimental stability limit taken from Fig. 32b.

8.6 Effect of pH and aqueous Ca concentration on the U(VI) and Np(VI)

sorption onto TiO2

The speciation calculations presented in Section 7.1 show that the Np(VI) and U(VI) speciation depends on the pH in the range 10 < pH < 14. In addition, the presence of Ca and Si in solution might influence the speciation of the hexavalent actinides in alkaline solutions, although the stabilities of ternary Ca-An(VI)-OH complexes and ternary U(VI)-hydroxy-silicate complexes have not yet been explored in detail. Sorption tests onto TiO2 were carried out at different pH values to evaluate the effect of pH and Ca concentration onto An(VI) surface complexation.

Np(VI) was found to sorb strongly on TiO2 at pH = 10 (Rd = 106 L kg−1) (Fig. 38a). Such a high Rd value was expected in the case of Np(VI) with a view to its effective charge (3.3), and the previously discussed correlation between ionic bond strength and effective charge (see Section 1.1). Furthermore, a high Rd value is in agreement with the sorption values determined for U(VI) on various solids under alkaline conditions in the absence of CO2 (e.g., Hsi and Langmuir, 1985; Bradbury and Baeyens, 2009). Fig. 38a further reveals that Np(VI) uptake by TiO2 decreased with increasing pH in the pH range 10 < pH < 14. The Np(VI) sorption data reported here and the U(VI) sorption data reported earlier by Comarmond et al. (2011) on the same TiO2 suggest that the Rd value for the sorption of hexavalent actinides reaches a maximum (∼ 106 L kg−1) in the pH range 7 < pH < 10 (Fig. 38a). This maximum Rd value coincides with the predominance of the anionic NpO2(OH)3

− or UO2(OH)3− species, respectively, at pH = 10

(Fig. 30c). The observed drop of the Rd value above pH = 10 corresponds very well with the increasing predominance of the anionic NpO2(OH)4

2− species in solution. These observations indicate that, notwithstanding its negative charge, the anionic species, NpO2(OH)3

-, strongly sorbs on the TiO2 surface. Hence, electrostatic repulsion does not prevent anionic Np(VI)-

10-10 10-9 10-8 10-7 10-6 10-5 10-410-6

10-5

10-4

10-3

10-2

10-1

100


Exp. 3 (233U) Exp. 2 (233U) Exp. 3 (Unat)

Ther

mod

ynam

ic U(

VI)

solu

bilit

y lim

it

Expe

rimen

tal U

(VI)

stab

ility

lim

itU(VI

) sorb

ed (

mol

es k

g-1 ) a

10-10 10-9 10-8 10-7 10-6 10-5 10-4101

102

103

104

105

Expe

rimen

tal U

(VI)

stab

ility

lim

it

Ther

mod

ynam

ic U(

VI)

solu

bilit

y lim

it

Exp. 3 (233U) Exp. 2 (233U) Exp. 3 (Unat)

R d(U(V

I)) (

L kg

-1)


b


hydroxy species from sorbing on negatively charged TiO2 surface sites in contrast to the anionic NpO2(OH)2

-.

Referring to the concept of electrostatic inter ligand repulsion (Section 2.3), the limiting number of OH groups that fit in the equatorial plane of the UO2

2+ and NpO22+ ion is nlimit = 4, which

implies that NpO2(OH)42- and UO2(OH)4

2- anions cannot further enhance their coordination sphere by binding to Ti-OH sites while binding is possible in the case of the NpO2(OH)3

- and UO2(OH)3

- species. A reduction factor, Fred, was defined based upon Eq. (3.14) (Section 3.4.2) to account for sorption reduction due to the formation of these anions:

0 * 'd 4

redd

R KF 1R [H ]+

= = + , (8.2)

with *K4’ being the conditional stepwise stability constant for the reaction:

22 3 2 2 4NpO (OH) H O NpO (OH) H− − ++ ⇔ + (8.3)

The Rd values in the pH range 10 < pH < 14 can then be calculated as Rd = Rd0/Fred. Log*K4

0 values at I = 0 are listed in Table 8. Ionic strength corrections for the calculation of the conditional stability constants, log*K4’, was based on the ionic strength corrections in accordance with the Davies equation (b = 0.3). The pH dependence of the Rd values for Np(VI) sorption on TiO2 could successfully be described with this model (Fig. 38a). This finding supports the idea that NpVIO2(OH)4

2- species is a limiting hydroxy complex which cannot form surface bonds with Ti-OH groups on TiO2. Note that the above reasoning implies that the same solid phase controls the sorption over the complete range of C:S ratios. In the case of the C-S-H phases it was shown that the number of sorption sites on the C–S–H phases is almost independent of the C:S ratio (Section 2.1.2.). On the assumption that the number of sorption sites is to a large degree responsible for the sorption properties of the C–S–H phases, we assume that the sorption properties of C–S–H phases are independent of the C:S ratio and thus that we can use the same Rd

0 value for all C:S ratios.

The effect of the aqueous Ca concentration on the Np(VI) uptake by TiO2 in 0.3 M NaOH at pH = 13.3 is shown in Fig. 38b. Np(VI) sorption was found to increase with increasing aqueous Ca concentration up to an equilibrium Ca concentration of ∼10−5 M.

Table 8: Overall and stepwise formation constants for relevant Np(VI) hydrolysis complexes used in the model calculations. (log*Kn

0 = log*βn0 - log*βn-1).

log*βn0 log*K4

0

UO2(OH)3-

UO2(OH)42-

-20.25a)

-32.4a)

-11.2

NpO2(OH)3-

NpO2(OH)42-

-21.2b)

-32.0b)

-10.8 a) Altmaier et al. (2017) b) Gaona et al., (2013a)

Ca is known to sorb onto the surface of TiO2 (Tits et al., 2014b). At equilibrium Ca concentration ∼10−5 M site saturation on TiO2 by Ca sorption is reached. Above this concentration, the sorbing Np(VI) thus binds to a “Ca-OH” like surface rather than to a “Ti-OH” like surface. The effect of Ca sorption on Np(VI) sorption can be explained in a similar way as


for Np(V), i.e. either by neutralization of the negative surface charge or by the formation of a very strong surface stabilized Ca-neptunate complex or Ca–neptunate ion pair. A thorough understanding of Np(VI) sorption processes in the presence of Ca is still lacking and would require more in-depth wet chemistry and spectroscopic investigations. At the time being we infer that both pH and Ca concentration have a pronounced effect on Np(VI) surface complexation on TiO2 in a way similar to Np(V) surface complexation. It should be noted that the same effects might be observed when Np(VI) sorbs onto C–S–H phases via a surface complexation process. Detailed studies are currently pending.

The effect of the aqueous silicate concentration and the formation of ternary An(VI)-hydroxy silicate complexes on the uptake of hexavalent actinides has not been investigated yet.

Fig. 38: Sorption of Np(VI) on TiO2. a) Effect of pH in absence of Ca. b) Effect of the Ca2+ concentration in 0.3 M NaOH. Experimental data (symbols) and modeling (line in Fig. a) (see text). Rd, max = 4.3·107 L kg-1. The line in Fig. 38b is added to guide the eye.

8.7 Effect of the C:S ratio

Sorption studies with U(VI) and Np(VI) were performed on C–S–H phases as function of the C:S ratio in alkali-free solutions in the pH range 10 < pH < 12.5 and in ACW at pH =13.3 with the aim of obtaining a more detailed understanding of the influence of the U(VI) and Np(VI) speciation on sorption (Fig. 39). In alkali-free conditions, the Rd values were very high at low C:S ratios, reaching values up to 106 L kg-1. With increasing C:S ratio, however, the Rd decreased steadily. In ACW, at constant pH = 13.3, the opposite trend was observed: Sorption onto C–S–H phases was relatively weak at low C:S ratios as indicated by the relatively low Rd values of ~600 L kg-1 in the case of Np(VI) and 2000 L kg-1 in the case of U(VI). With increasing C:S ratio, sorption increased with Rd reaching values of 2500 L kg-1 (Np(VI)) and 4·104 L kg-1 (U(VI)). Generally, the sorption of U(VI) tended to be stronger than the sorption of Np(VI). This observation is consistent with the trends observed in hydrolysis, i.e. a higher stability of the U(VI) hydroxy complexes compared to the stability of the corresponding Np(VI) complexes.

2 4 6 8 10 12 14101

102

103

104

105

106

107

Np(VI): this workU(VI): Comarmond et al. 2011

R d (L

kg-1

)

pHc

a

Slope = -1

10-6 10-5 10-4 10-30102

103

104

105

R d (N

p(VI

) (L

kg-1

)

Ca2+ concentration (M)

0.3 M NaOHb


Fig. 39: Uptake of U(VI) and Np(VI) by C–S–H phases under alkali-free conditions (10 < pH < 12.5) and in ACW at a fixed pH of 13.3. a) Effect of the C:S ratio, b) Rd values from a) plotted versus pHc, c) Rd values determined in ACW at a constant pH = 13.3, and plotted versus the aqueous Ca2+ concentration. Rd, max = 4.3·107 L kg-1 (239Np experiment) and 1.5·107 L kg-1 (233U experiment).

The Np(VI) sorption data on C–S–H phases were plotted as function of pH and aqueous Ca concentration to assess the influence of the aqueous speciation on the sorption behaviour. A reversible adsorption process, similar to that on TiO2, would imply that with increasing C:S ratio, the Rd values were affected simultaneously by two counteracting parameters: 1) A decrease in the Rd value with increasing pH due to increasing concentrations of the non-sorbing anionic UO2(OH)4

2- or NpO2(OH)42− species in solution, and 2) an increase in Rd value with

increasing aqueous Ca concentration. Both effects may partially compensate each other, which limits a direct comparison of the sorption behaviour on both solids. The sole effect of the aqueous Ca concentration on Np uptake by C–S–H phases is apparent from the sorption tests in the presence of alkalis (ACW, pH = 13.3) (Fig. 39c). An increasing C:S ratio which corresponds to an increasing aqueous Ca concentration, appears to promote the sorption of anionic UO2(OH)4

2- or NpO2(OH)42− species onto C–S–H phases. The Rd value for Np(VI) rises from

6·102 L kg−1 at a C:S ratio of 0.67 (corresponding to an aqueous Ca concentration of 2·10−5 M) to 2.5·103 L kg−1 at a C:S ratio of 1.25 (corresponding to an aqueous Ca concentration of 1.6·10−3 M). This indicates that an increase in the aqueous Ca concentration of about two orders of magnitude gives rise to an increase in the Rd value by a factor 4.

The sorption data presented in Fig. 39b were modelled using the reduction factor Fred defined in Eq (3.14) and setting the value for Rd

0 to the Rd value at pH 10 with the aim of determining the

0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0101

102

103

104

105

106

107

108

Rd,max (Np(VI))

Np(VI) Alkali-free Np(VI) in ACW U(VI) Alkali free U(VI) in ACW

R d (L

kg-1

)

C:S (mol/mol)

Rd,max (U(VI))

a

10 11 12 13 14102

103

104

105

106

107

108

Rd,max(Np(VI))

Np(VI) Alkali-free Np(VI) ACW U(VI) Alkali-free U(VI) ACW

R d (

L kg

-1)

pHc

b

Rd,max(U(VI))

10-5 10-4 10-3101

102

103

104

105

106

107

108Rd,max(Np(VI))

U(VI) Np(VI)

R d (

L kg

-1)

Ca2+ equilibrium concentration (M)

ACW at pH = 13.3

cRd,max(U(VI))


contribution of the aqueous Np(VI) speciation to its sorption behaviour on C–S–H phases. The solid line in Fig. 39b shows that this model approach explains quite well the pH dependence of the sorption data. The effect of the C:S ratio on the sorption of Np(VI) and U(VI) by C–S–H phases can be attributed to the formation of the non-sorbing NpO2(OH)4

2- and UO2(OH)42-

complexes, respectively. The enhanced sorption of U(VI) compared to Np(VI) is thoroughly explained by the higher stepwise stability constant for Np(VI) (log*K4

0 = -10.8) compared to U(VI) (log*K4

0 = -12.15). There could be a small effect of the Ca concentration on sorption as indicated in Fig. 39c. But this effect appears to be negligible compared to the hydrolysis effect.

8.8 Discussion

8.8.1 Sorption data The sorption behaviour of An(VI), in particular U(VI) and Np(VI), on C–S–H phases and HCP can be summarized as follows:

1) Strong sorption illustrated by Rd values in the range of ~5·102 L kg-1 < Rd < ~3·106 L kg-1;

2) Relatively fast sorption kinetics onto C–S–H phases (steady state attained in ~3 days), and apparently slower sorption kinetics onto HCP (~30 days);

3) Non-linear sorption onto C–S–H phases in alkali-free conditions and in ACW, and linear sorption onto HCP;

4) Decreasing Np(VI) sorption with increasing pH on TiO2 due to the formation of the non-sorbing NpO2(OH)4

2- species;

5) Increasing sorption on TiO2 with increasing Ca2+ concentration in solution;

6) Decreasing Rd value with increasing C:S ratio of C–S–H phases in alkali-free conditions largely explained by hydrolysis effects on the assumption that the NpO2(OH)4

2- and UO2(OH)4

2- species are not sorbing. Increasing Rd values with increasing C:S ratio in ACW at a constant pH of 13.3 explained by the positive effect of Ca on An(VI) sorption.

Several studies on the uptake of hexavalent actinides (mainly U(VI)) have been published in the literature in the past (Sutton et al., 2003; Pointeau et al., 2004a; 2004b; Tits et al., 2011). Comprehensive reviews of the existing sorption data are available from Ochs et al. (2016) and Wieland (2014). Ochs et al. (2016) compiled sorption values for U(VI) sorption onto various cement types in a wide range of chemical conditions (10 < pH < 12.5) and degradation states, published in British scientific reports. These Rd values appear to be relatively independent of salinity, pH, degradation stage and temperature, and range between ~2·103 L kg-and 104 L kg-1. Pointeau et al. (Pointeau et al., 2004b; Pointeau et al., 2008) carried out very detailed studies of U(VI) sorption on C–S–H phases with various C:S ratios and HCP. The authors observed linear and reversible sorption in contrast to the present study but this apparent inconsistency might be attributed to the high scatter of the data and the lack of experimental data in the higher concentration range in the study of Pointeau et al. (2004b). The same authors also observed an increase in Rd value with decreasing pH (progressing cement degradation) which they attributed to a decreasing competition with Ca for sorption sites. The sorption data presented in this study (Fig. 39c) contradict this explanation because Ca clearly has a positive effect on U(VI) and


Np(VI) sorption at constant pH. The alternative explanation for the pH effect is hydrolysis, i.e. the decreasing concentration of the non-sorbing UO2(OH)4

2- and NpO2 (OH)42- species with

decreasing pH.

8.8.2 Uptake mechanisms Several spectroscopic studies were carried out to obtain information on the local coordination environment of U(VI) sorbed onto cement minerals. XAS investigations were not conclusive due to difficulties encountered with the interpretation of the EXAFS spectra. In most spectra, backscattering contributions of neighbouring cations from the C–S–H phases were very weak due to suppression by strong backscattering paths of the linear actinyl moieties (Harfouche et al., 2006; Macé et al., 2013). Nevertheless, these studies revealed that U(VI) is sorbed onto C–S–H phases in two different coordination environments depending on the C:S ratio of the C–S–H phases (Gaona et al., 2012b) and that these coordination environments show some similarities with those of uranyl silicates, e.g. a split equatorial oxygen shell, neighbouring Si atoms at short and long distances, and Ca neighbours (Harfouche et al., 2006). Incorporation into the main Ca-O layer is not very probable because participation of Ca in the coordination sphere of U(VI) or Np(VI) was not observed in the EXAFS studies or Ca was detected only at long distances (3.9Å – 4.3 Å) from the central U(VI) atoms (Harfouche et al., 2006), respectively. At high U(VI) loading a Ca-uranate-like precipitate was shown to be formed (Macé et al., 2013). Harfouche et al. (2006) proposed a structure in which a sorbed uranyl moiety is located in the C–S–H interlayer, bound in a bidentate fashion to a bridging Si-tetrahedron of one silica chain and in a monodentate fashion to a Si-tetrahedron of a second silica chain. This coordination environment could, however, not be confirmed in the XAS studies with Np(VI) carried out by Gaona et al. (2013b).

Tits and co-workers investigated the uptake of U(VI) by C–S–H phases with different C:S ratios and increasing U(VI) loadings using luminescence spectroscopy (Tits et al., 2011; 2015). Their findings are largely in agreement with those of the XAS investigations. Three sorbed U(VI) species were detected, a U(VI) surface complex, a U(VI) sorbed species located in the C–S–H interlayer and, at high loading a Ca-uranate-like surface precipitate. U(VI) solid speciation was found to change continuously with time over a period up to 6 months notwithstanding the fast sorption kinetics observed in batch sorption studies. This discrepancy may be explained by the large uncertainties on the Rd values measured in the batch sorption studies, thus covering the slight increase in sorption with time due to the progressing relocation of U(VI) from the C–S–H surfaces into the interlayer. The presence of two different U(VI) sorbed species is further in agreement with the non-linear sorption behaviour observed in the sorption isotherm measurements (Section 8.4).

Gaona et al. (2012a) proposed a geochemical model of the U(VI) uptake by C–S–H phases based upon sublattices occupied with U(VI) bearing species in accordance with information obtained from spectroscopic studies. The proposed solid solution model contained six end-members and was found to satisfactorily describe U(VI) sorption behaviour on cementitious materials found in the literature.

The effect of [Ca] on the U(VI) and Np(VI) sorption onto C–S–H phases and also onto TiO2 is evident but not yet understood in detail. A further gap in our understanding of the An(VI) behaviour in cementitious environments concerns their interaction with aqueous silica and the formation and stability of aqueous ternary An(VI)-hydroxy silicate complexes. These complexes could be important at high silica concentrations and low pH, such as in low-pH cements.


9 Summary and conclusions

The present report summarizes the batch sorption studies on the uptake of actinides by cementitious materials carried out at LES/PSI in support of the safety assessment. In these studies, the sorption of trivalent, tetravalent, pentavalent and hexavalent actinides onto C–S–H phases and HCP has been investigated. For some actinides batch sorption experiments were complemented with co-precipitation experiments under highly oversaturated conditions to explore the role of structural incorporation as an actinide sorption mechanism in C–S–H phases. Finally, actinide sorption values onto C–S–H phases were compared to sorption values onto HCP to validate the starting assumption that C–S–H phases are the uptake controlling cement component for actinides. The latter assumption was developed in the framework of spectroscopic investigations carried out at LES/PSI on actinide/lanthanide interaction with cementitious materials.

Interaction of tri-, tetra, penta, and hexavalent actinides with cementitious materials is very strong, thus leading to very high Rd values (typically 103 L kg-1 > Rd > 106 L kg-1) in the chemical conditions of cement systems. Hydrolysis of the actinides has a major effect on the extent of actinide interaction with the cementitious materials. The effect of hydrolysis and therefore the affinity of the different actinides for C–S–H phases and HCP can be rationalized in terms of the concept of electrostatic inter-ligand repulsion predicting that, for each oxidation state of a given actinide, only a limiting number of OH groups, nlimit, can fit in its first coordination sphere. The nlimit values reported in the literature for the actinides under investigation are 4 (An(III)), > 4 (An(IV)), 2 (An(V)) and 4 (An(VI)), meaning that the actinide species AnIII(OH)4

-, AnVO2(OH)2- and AnVI(OH)4

2- cannot make surface complexes with silanol groups or silandiol groups of C–S–H phases. With this information at hand, a sorption reduction factor, Fred, was defined corresponding to the ratio of the Rd value in the absence (Rd

0) and in the presence (Rd) of the non-sorbing actinide hydroxy species. This concept has proven useful to understand pH effects on sorption while it ignores the presence of different sorption sites with different affinities for radionuclide binding in C-S-H phases. Application of the conceptual Fred approach allowed to satisfactorily describe the influence of the C:S ratio on the sorption of hexavalent actinides onto C–S–H phases and showed that the sorption of these actinides could almost entirely be controlled by their hydrolysis.

The batch sorption studies on C–S–H phases reveal a sorption behaviour typical for a pure surface-controlled process, i.e., fast sorption reactions, in most cases (quasi-) linear sorption, sorption dependence on aqueous speciation and, last but not least, reversible sorption processes. At first glance, this macroscopic sorption behaviour appears to be in contradiction with the large body of spectroscopic evidence for slow structural incorporation processes reported in the literature. However, the large uncertainties inherent to batch sorption experiments with actinides on C–S–H phases may actually cover slow incorporation kinetics and sorption – desorption hysteresis. The lack of hysteresis (similar sorption and desorption kinetics) of the actinide / lanthanide sorption suggest that the incorporation processes are rather fast processes as well. The results from the present batch sorption studies suggest that the slow incorporation process following rapid surface complexation does not exert a significant influence on the macroscopic sorption behaviour (Rd value, influence of the aqueous speciation, sorption hysteresis) of the actinides onto C–S–H phases and that this sorption behaviour can macroscopically be described by surface sorption models.

Sorption tests with Np(V) and Np(VI) onto TiO2 further indicated a positive effect of the aqueous Ca concentration on surface complexation of these actinides in alkaline conditions. Increasing Ca concentrations had a significant influence on the Np(V) sorption onto C–S–H phases in that it completely neutralized the negative effect of the progressing dominance of the


non-sorbing NpO2(OH)2- species with increasing C:S ratio. Therefore, the experimentally

observed uptake of Np(V) by C–S–H phases was much stronger than predictions made earlier on the basis of the effective charge of Np(V). The influence of Ca on the sorption of An(VI) onto C–S–H phases was found to be marginal. The mechanism responsible the positive effect of Ca on actinides sorption in alkaline conditions is not yet understood in detail.

As a concluding remark we would like to draw attention to still existing gaps in our understanding of the actinide speciation in solution. The present study demonstrates that a good knowledge of the aqueous actinide speciation in alkaline cementitious pore waters is a prerequisite for predicting the actinide sorption behaviour in the cementitious near-field of a nuclear waste repository. While the knowledge on the formation of ternary Ca-AN-OH complexes has significantly improved during the last decade, an important lack of knowledge still exists on the actinide complexation with aqueous silica and the formation and stability of aqueous ternary An-hydroxyl-silicate complexes. The existence and stability of such complexes has already been reported in the literature in the case of Th(IV). It is expected that similar complexes exist for the other actinides. An-hydroxyl-silicate complexes may be important species in more complex cementitious systems, such as low-pH cements, when Portland cement is mixed with supplementary cementitious materials such as silica fume, fly ash and blast furnace slag.

10 Acknowledgements

The authors would like to thank A. Schaible (PSI), J. P. Dobler (PSI), D. Kunz (PSI), A. Laube (PSI) and M. Mantovani (PSI) for their assistance with the experiments. Two reviewers, Dr. X. Gaona and Dr. M. Felipe-Sotelo, are acknowledged for their detailed reviews. Their remarks and corrections have significantly improved the quality of this report. Partial financial support by the Swiss National Cooperative for the Disposal of Radioactive Waste (Nagra) is kindly acknowledged.

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