February 2014

Working Reports contain information on work in progress

or pending completion.

Arto Muurinen

VTT

Christophe Tournassat, Jebri l Hadi

BRGM

Jean-Marc Greneche

LPCE

Working Report 2014-04

Sorption and Diffusion of Fe(II) in Bentonite

 

ABSTRACT  The iron in the engineering barrier system of a nuclear waste repository interacts via the corrosion process with the swelling clay intended as the buffer material. This interaction may affect the sealing properties of the clay. In the case of iron-bentonite interaction, redox reactions, dissolution/precipitation, the diffusion and sorption are coupled together. In a combined study different processes are difficult to distinguish from each other, and more specific studies are needed for the separate processes. In particular, there is a need for well-controlled diffusion and sorption experiments where iron is kept as Fe(II). In this project, sorption and diffusion of Fe(II) in bentonite have been studied. The experiments were carried out under low-oxygen conditions in an anaerobic glove-box. The radioactive isotope (55Fe) was used as a tracer in the experiments. The sorption experiments were carried out with two batches of purified MX-80 bentonite. One was purified at Bureau de Recherches Géologiques et Minières, French Geological Survey (BRGM) and the other one at VTT Technical Research Centre of Finland (VTT). Experiments were also carried out with synthetic smectite, which did not include iron, which was prepared at LMPC (ENSC, F 68093 Mulhouse, France). The sorption experiments were carried out in 0.3 M and 0.05 M NaCl solutions as a function of pH, and in 0.3 M NaCl solution buffered at pH 5 as a function of added Fe(II) concentration. The separation of bentonite and solution at the end of the sorption experiment was carried out in the early phase by centrifuging only. In the later phase, ultrafiltering was added in order to improve the separation. The diffusion experiments were carried out in compacted samples prepared from MX-80 purified at VTT and saturated with 0.3 M NaCl at pH 8 and 5. A non-steady-state diffusion experiment method, where the tracer was introduced as an impulse source between two bentonite plugs was used in the measurements. Qualitatively, the sorption of the radioactive 55Fe on all the clays shows the same type of behaviour, i.e. sorption increases with increasing pH. In the measurements with the bentonite purified at VTT, the sorption occurs at a higher pH than in the measurements carried out with bentonite purified at BRGM. The sorption experiments in the acetate buffer of pH 5 show decreasing sorption of 55Fe as a function of the increasing concentration of the added Fe(II). A general model for the investigated clays is proposed where Fe sorption is due to adsorption on exchange sites, strong and weak complexation sites and electron transfer with the structural Fe. All mechanisms identified apply to all clay samples but with variations in CEC values, structural Fe redox potential and strong and weak sites' surface density. The measured diffusivities show rather low values (10-15 – 10-16 m2/s) at pH 8 and 5. At pH 8, the diffusion curve calculated with a reactive transport model on the basis of the sorption matches fairly well the experimental results. At pH 5, the model predicts a much longer diffusion distance than found in the experiment. The reason for this discrepancy is not yet understood. A possible explanation could be a slow redox/sorption process which does not appear in the short batch sorption measurements. Keywords: Bentonite, Fe(II), sorption, diffusion.

FE(II):N DIFFUUSIO JA SORPTIO BENTONIITISSA TIIVISTELMÄ Ydinjätteen loppusijoitustilassa oleva rauta vuorovaikuttaa korroosioprosessin kautta puskurimateriaaliksi aiotun paisuvan saven kanssa. Tällä vuorovaikutuksella voi olla vaikutusta saven tiivistysominaisuuksiin. Rauta-bentoniitti vuorovaikutuksen tapauk-sessa redox-reaktiot, liukeneminen/saostuminen, diffuusio ja sorptio ovat kytkeytyneenä yhteen. Kytketyssä tutkimuksessa eri prosessit on vaikea erottaa toisistaan ja täs-mällisempiä tutkimuksia tarvitaan erillisille prosesseille. Tarvetta on varsinkin hyvin kontrolloiduille diffuusio- ja sorptiokokeille, joissa rauta on Fe(II):na. Tässä projektissa tutkittiin Fe(II):n sorptiota bentoniittiin. Kokeet tehtiin alhaisessa happipitoisuudessa n.s. hapettomassa hansikaskaapissa käyttämällä radioaktiivista 55Fe isotooppia merkkiaineena. Kokeita tehtiin kahdella erällä puhdistettua MX-80 bento-niittia. Toinen erä oli puhdistettu BRGM:ssä (Bureau de Recherche Géologique et Minière, French Geological Survey) ja toinen VTT:llä (Teknologian tutkimuskeskus VTT). Kokeita tehtiin myös LMPC:ssa (ENSC, F 68093 Mulhouse, France) valmis-tetulla synteettisellä smektiitillä, joka ei sisältänyt rautaa. Sorptiokokeet tehtiin 0.3 M ja 0.05 M NaCl-liuoksessa pH:n funktiona ja pH 5:teen puskuroidussa 0.3 M NaCl-liuoksessa lisätyn Fe(II)-pitoisuuden funktiona. Bentoniitin ja liuoksen erottaminen tehtiin alkuvaiheessa vain sentrifugoimalla. Myöhemmässä vaiheessa lisättiin ultra-suodatus erotuksen parantamiseksi. Diffuusiokokeet tehtiin puristetuilla näytteillä, jotka oli valmistettu VTT:llä puhdistetusta MX-80 bentoniitista ja kyllästetty 0.3 M NaCl-liuoksella pH:ssa 8 ja 5. Diffuusiokokeet tehtiin menetelmällä, jossa merkkiaine lisättiin impulssilähteenä kahden bentoniittipalan väliin. Kvalitatiivisesti radioaktiivisen 55Fe:n sorptio on kaikilla savilla samantyyppistä, eli sorptio kasvaa pH:n noustessa. VTT:llä puhdistetulla bentoniitilla sorptio tapahtuu ylemmässä pH:ssa kuin BRGM:ssä puhdistetulla bentoniitilla. Fe-55:n sorptio ase-taattipusturissa, jonka pH oli 5 pieneni liuokseen lisätyn Fe(II):n pitoisuuden kasvaessa. Tutkituille saville ehdotetaan yleistä mallia, jossa raudan sorptio aiheutuu adsorptiosta ioninvaihtopaikkoihin, vahvoihin ja heikkoihin kompleksaatio paikkoihin ja elektronin siirrosta rakenteellisen raudan kanssa. Kaikki tunnistetut mekanismit sopivat kaikkiin savinäytteisiin kuitenkin erilaisella CEC arvolla, rakenteellisen raudan redoxpo-tentiaalilla ja vahvan ja heikon pintapaikan tiheydellä. Mitatut diffuusiokertoimen arvot olivat alhaisia (10-15 – 10-16 m2/s) pH:ssa 8 ja 5. Reaktiivisella kulkeutumismallilla sorptiotuloksen pohjalta laskettu diffuusiokäyrä vastasi pH:ssa 8 suhteellisen hyvin mittaustulosta. Kuitenkin pH:ssa 5 malli ehdottaa paljon pidempää diffuusiomatkaa kuin mittauksessa havaittiin. Syytä tähän eroon ei vielä ymmärretä. Mahdollisena selittäjänä voisi olla hidas redox-/sorptioprosessi, joka ei ilmene lyhyessä sorptiomittauksessa. Avainsanat: Bentoniitti, Fe(II), sorptio, diffuusio.

PREFACE Arto Muurinen (VTT, Finland) was responsible for the sorption and diffusion experiments of this study. The modelling work was performed by Christophe Tournassat (BRGM, France). Jean-Marc Greneche (LPCE, Le Mans, France) and Jebril Hadi (BRGM / ISTERRE) provided Mössbauer spectrometry results. Jocelyne Brendlé (LMPC, France) is thanked for having provided Fe-free synthetic clay samples. The contact person at Posiva was Marjut Vähänen.

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TABLE OF CONTENTS ABSTRACT TIIVISTELMÄ PREFACE 1  INTRODUCTION .................................................................................................... 3  2  EXPERIMENTAL .................................................................................................... 5 

2.1  Sorption experiments ....................................................................................... 5 2.1.1  Bentonites used in the study ..................................................................... 5 2.1.2  Arrangements used in the sorption experiments ....................................... 8 2.1.3  Sorption experiments in 0.3 M NaCl at varying pH ................................. 10 2.1.4  Sorption experiments in 0.05 M NaCl at varying pH ............................... 10 2.1.5  Sorption experiments in 0.3 M NaCl solution at constant pH 5 in varying Fe concentrations .................................................................................... 11 2.1.6  Results of the sorption experiments ........................................................ 11 

2.2  Diffusion experiments ..................................................................................... 25 2.2.1  Preparation of the samples for the diffusion experiments at pH 8 .......... 26 2.2.2  Preparation of samples for the diffusion experiments at pH 5 ................ 26 2.2.3  Storing and finishing of the diffusion experiments ................................... 28 2.2.4  Comparison of measured and calculated diffusion coefficients .............. 30 2.2.5  Measurement of pH in bentonite samples ............................................... 31 

3  MODELLING OF THE EXPERIMENTS ................................................................ 35 

3.1  Data selection and corrections ....................................................................... 35 3.1.1  Problem of solid / liquid separation ......................................................... 35 3.1.2  Desorption vs. dissolution at low pH ....................................................... 37 

3.2  General considerations for modelling ............................................................. 41 3.3  Modelling Fe sorption on synthetic smectite with a cation exchange / surface complexation model ....................................................................................... 42 3.4  Modelling Fe sorption on synthetic smectite with a cation exchange / surface complexation model, taking into account surface redox reactions ................. 45 3.5  Modelling structural iron redox properties in natural clays ............................. 49 3.6  Modelling of Fe(II) sorption on natural montmorillonite .................................. 54 

3.6.1  MXBRGM .................................................................................................... 54 3.6.2  MXVTT ...................................................................................................... 56 3.6.3  from Ph.D. thesis of Tournassat (2003) .................................................. 57 3.6.4  Summary of modelling results and challenges in finding a general model for Fe(II) sorption on natural montmorillonites ........................................ 57 3.6.5  Proposition of a general model for montmorillonites and other smectites ... ................................................................................................................ 57 

3.7  Application of other data from the literature ................................................... 64 3.8  Conclusion on the applicability of the model .................................................. 65 3.9  Modelling of the diffusion experiments ........................................................... 65 

4  SUMMARY ........................................................................................................... 69  REFERENCES ............................................................................................................. 71  APPENDICES ............................................................................................................... 75

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1 INTRODUCTION The iron in the EBS of a spent fuel repository interacts via the corrosion process with the swelling clay that is intended as the buffer material. This interaction may, in principle, affect the sealing properties of the clay. In the case of iron-bentonite interaction, redox reactions, precipitation, diffusion and sorption are coupled together. In a combined study like that by Carlson et al. (2007), different processes are difficult to distinguish from each other, and more specific studies are needed for the separate processes. The diffusion properties of Fe2+ has been considered one of the topic that requires further studies in Wersin and Snellman (2008). Diffusion is one of the essential processes that determine how far the effects of the corroding iron can enter into the clay. In the sorption and diffusion studies, the concentrations are typically kept low in order to avoid precipitation. In the case of Fe2+ oxidation, the solubility can decrease dramatically and lead to precipitation of Fe3+. Because the iron and the radionuclides may adsorb on the same sites of the bentonite, there is competition over the sorption sites. The sorption of radionuclides could then be reduced owing to the occupation of the sorption sites by other species like Fe2+. Indeed, Fe(II) – Zn(II) competitive sorption experiments clearly show that Fe(II) displaces Zn(II) on clay edge sites (Tournassat, 2003). Very few diffusion experiments have been carried out with iron in bentonite. Idemitsu et al. (1993, 2002) used 59Fe as a tracer to measure the apparent diffusivities of iron with and without metallic iron in contact with the bentonite. The apparent diffusivity was about 10-12 m2/s when metallic iron was present and about 10-14 m2/s when metallic iron was not present. It was concluded that the diffusing species were Fe2+ in the former case and Fe3+ in the latter case. Sorption is a process that can decrease the apparent diffusivities in a porous medium. Two types of sorption processes are usually expected to occur in bentonite, i.e. cation exchange and surface complexation. Tournassat et al. (2009) concluded that Fe2+ has almost the same behaviour and affinity as Ca2+ on a clay exchanger site. Bradbury and Baeyens have published many reports and papers on the characterization of the cation exchange, surface complexation and protonation/deprotonation reactions of bentonite. Systematic sorption measurements and modelling of the results have been performed with Ni and Zn in bentonite (Baeyens and Bradbury 1997, Bradbury and Baeyens 1997). A linear free energy relationship was used in Bradbury and Baeyens (2005) to estimate the surface complexation constants for Mn(II), Ni(II), Zn(II), Cd(II), Eu(III), Am(III), Sn(IV), Th(IV), Np(V) and U(VI) on the basis of in-house and literature sorption data. Hunter et al. (2007) used the equations presented by Bradbury and Baeyens (2005) for the linear free energy relationship to evaluate the surface complexation constants for Fe2+. They have also presented the cation-exchange constant according to Gaines’ and Thomas’ convention on the basis of the work by Kamei et al. (1999). The evaluated constants for different reactions are presented in Table 1.

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Table 1. Surface complexation and cation-exchange constants for Fe2+ in bentonite determined with the linear free energy relationship (Hunter et al. 2007).

Strong site surface complexation for iron log sKx-1 ≡Surfa_strongOH + Fe2+ ↔ ≡Surfa_strongOFe+ + H+ -0.45 ≡Surfa_strongOH + Fe2+ + H2O ↔ ≡Surfa_strongOFeOH + 2H+ -10.4 ≡Surfa_strongOH + Fe2+ + 2H2O ↔ ≡Surfa_strongOFe(OH)2

- + 3H+ -19.8 Weak site surface complexation for iron log wKx-1 ≡Surfa_weakoneOH + Fe2+ ↔ ≡Surfa_weakoneOFe+ H+ -3.11 Cation exchange for iron log K 2 Na-montmorillonite + Fe2+ ↔ Fe-montmorillonite + 2 Na+ 0.267 Moreover, if modelling Fe retention data from Tournassat (2003) on natural montmorillonite by surface complexation only, the complexation constants for Fe(II)-specific sorption on clay edge sites are much higher than those published for other divalent cations and also much higher than the value estimated from the LFER estimate of Hunter et al. (2007). It is important then to assess the exact nature and extent of Fe(II) sorption on natural montmorillonite. According to Tournassat (2003), Fe2+ sorption in bentonite is influenced by several different reactions like cation exchange on the surfaces, specific pH-dependent sorption on the edge sites, surface sorption together with H4SiO4 and a sorption followed by an oxidation reaction. The relative importance of different sorption mechanisms on montmorillonite edge surfaces appeared to be dependent on the pH. In the alkaline pH range, an Si-Fe surface precipitate forms if a sufficient concentration of H4SiO4 is available. In the near-neutral pH, the high Fe(II) sorption is explained by, at least, a two-step mechanism including an adsorption step followed by an oxidation step. All of these sorption mechanisms were studied at relatively high Fe(II) concentration (above 1 µmol L-1) because (i) they were carried out with natural Fe(II) (as a consequence it was not possible to measure Fe(II) at very low concentration using conventional methods), and (ii) natural montmorillonite itself contains Fe(II), which dissolves in the course of the experiment and leads to an unavoidable background Fe(II) concentration causing the error band on the Fe(II) sorption estimation. Point (ii) was further addressed by Gehin et al. (2007) using a synthetic montmorillonite free of Fe in its structure. While the two-step uptake mechanism (adsorption followed by partial oxidation) has been identified again in this system, the Fe(II) surface complexation constant was lower than that encountered for natural montmorillonite and more in line with the value estimated by Hunter (2007). To our knowledge, point (i) has not been addressed by using an Fe radiotracer, although it would be essential for quantifying and modelling Fe(II) sorption in natural conditions (at pH 7–8, a 100 % value is found owing to the measurement imprecision).

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2 EXPERIMENTAL 2.1 Sorption experiments 2.1.1 Bentonites used in the study The sorption tests with Fe(II) were carried out with two batches of purified MX-80 bentonite, and a synthetic smectite. One of the MX-80 bentonites was purified by BRGM (MXBRGM, Tournassat, 2003), and the other one by VTT (MXVTT). The synthetic smectite was obtained from BRGM (SBBRGM, Gehin et al., 2007). The method described in Tributh and Lagaly (1986) was followed to purify the MX-80 bentonite at VTT. The purification comprised removal of the large particles, dissolution of carbonates by acid treatment, dissolution of iron oxides, removal of organic material, changing to the sodium form, and removal of the excess salt by washing and finally by dialysis. Removal of large particles Twelve 35-gram batches of air-dry MX-80 bentonite were dispersed with ultrasound into 800 ml of deionized water. The batches were shaken overnight and centrifuged with a centrifugal field of 60 x g for nine minutes. The supernatant was collected in a large vessel. The bottom was dispersed twice in water and treated as above. Finally, the bottoms were thrown away. Sodium chloride was then added to the bentonite dispersion so that the concentration was 0.5 mol/l. The bentonite was separated from the solution by centrifuging in one-litre centrifuge bottles with a centrifugal field of 7000 x g. The final volume was about 5 litres. Dissolution of carbonates Five litres of solution containing 164 g/l sodium acetate and 120 g/l acetic acid was added into the five-litre batch of the bentonite dispersion. The batch was allowed to react for two and half days by mixing occasionally. The bentonite was then separated by centrifuging into six centrifuge bottles. The solution was thrown away. Removal of iron oxides To each of the six centrifuge bottles, 650 ml of citrate buffer was added. (Citrate buffer: 115 g Na-citratedihydrate, 8.5 g NaHCO3 and 70 g NaCl per litre). Bentonite was mixed into the solution. To each bottle, 20 g Na-dithionite was added and the mixture was heated to 70 oC. The dispersion was allowed to cool by occasionally shaking. The bottles were left to stand overnight. The next day the solution was separated in a centrifuge and removed. The bentonite was washed twice with 650 ml of solution comprising 0.5 mol/l NaCl and 0.01 mol/l HCl. Then the bentonites were washed twice with 0.5 M NaCl. Removal of organic material To each bottle, 500 ml of 0.1 M sodium acetate was added and then 150 ml 30-% hydrogen peroxide in small portions as quickly as the bubbling allowed. Then the mixture was heated slowly for four hours to 70 oC by mixing the slurry in order to decrease bubbling. Then it was allowed to cool and stand overnight. The next day the bentonite was separated in the centrifuge and washed three times with 1 M NaCl.

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Removal of excess salt The bentonite from the treatment with 1-molar NaCl was washed with deionized water until the bentonite could no longer be separated from the water. Then the bentonite dispersion was moved to dialysis tubes and dialysed until the electric conductivity in the solution was about 0.004 mS/cm. Drying and milling The bentonite was moved from the dialysis bags on plastic plates. Most of the water was removed from the bentonite by drying at 60oC. The rest of the water was moved into a vacuum freeze dryer. The dry bentonite was milled with an agate swing mill to a particle size < 1 mm. Preparation of bentonite dispersions Before use in the sorption experiments, the dry montmorillonites were stored for several weeks in a nitrogen glove-box in open vessels in order to remove the oxygen. Then dispersions of about 10 g/l were prepared from the montmorillonites and deionized water. The dispersions were stored for at least a week in a nitrogen glove-box and mixed with a magnetic bar in order to remove possible traces of oxygen. The volume was adjusted so that the final clay concentration was 10 g/l. Characterization of the CEC The cation-exchange capacity of the bentonite purified at BRGM and VTT and the CEC of the synthetic smectite prepared at BRGM were determined using the CuTrien method (Amman et al. 2005). Two parallel samples were used in the measurements. Table 2 presents the obtained results.

Table 2. Measured cation-exchange capacities (CEC) of MX-80 purified at BRGM and VTT and the CEC of the synthetic smectite prepared at BRGM.

Initial bentonite Processing CEC (meq/g) MX-80 purified at BRGM 0.791 MX-80 purified at VTT 0.975 Synthetic smectite prepared at BRGM 0.867 Measurement of Fe(II)/Fe(III) ratio in bentonite using Mössbauer spectrometry Smectite 57Fe Mössbauer spectra were recorded at 77 K using a constant acceleration spectrometer and a 57Co source diffused into a rhodium matrix. Velocity calibrations were made using -Fe foil at 300 K. The hyperfine parameters were refined using a least-squared fitting procedure (MOSFIT programme, Teillet and Varret, unpublished programme). The spectra for MXVTT and MXBRGM are shown in Figure 1. No spectra were measured for synthetic smectite as it contains no Fe in its structure.

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-12 -9 -6 -3 0 3 6 9 12

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Figure 1. 77K 57Fe Mössbauer spectra for MXVTT (left) and MXBRGM (right). Coloured lines corresponds to hyperfine contributions tabulated in Table 3.

Table 3. Mössbauer parameters for the samples MXVTT and MXBRGM.

Sample Contribution I.S. Q.S. or 2 Hhyp % -12 – 12 mm/s spectra (upper figures)

MX80_BRGM Fe(II) 1.33 0.72 3.05 7 Fe(III) 0.47 0.76 0.62 62 Magnetic sextet 0.35 0.8 0.38 55.1 6 Magnetic sextet 0.45 0.8 0 50.7 5 Magnetic sextet 0.76 0.8 0 43.7 5 Singlet 0.46 4 15 MX80_VTT Fe(II) 1.26 0.35 3.08 14 Fe(III) 0.45 0.88 0.67 52 Magnetic sextet 0.51 0.8 0 54.9 8 Magnetic sextet 0.72 0.8 0 44.6 5 Singlet 0.46 7 22

-3 – 3 mm/s spectra (lower figures)MX80_BRGM Fe(II) 1.28 0.6 3.0 7 Fe(III) 0.47 0.73 0.63 72 Singlet 0.45 3 21 MX80_VTT Fe(II) 1.26 0.38 3.08 19 Fe(III) 0.45 0.88 0.66 64 Singlet 0.45 5 17

I.S. = isomer shift; = line width; Q.S. = quadrupolar splitting value; 2= quadrupolar shift value, Hhyp = hyperfine field value, % = ratio of each component.

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Mössbauer spectra obtained from -12 to 12 mm/s show the presence of a magnetic contribution for both MXBRGM and MXVTT. A single line contribution is also present. Magnetic sextet is attributed to the presence of remaining iron oxides despite clay preparation procedure. The single line is probably due to the presence of superparamagnetic Fe(III)-containing particles, resulting from an incomplete chemical reaction. Quadrupolar contributions are attributed to Fe(II) and Fe(III) present in the structure of the clay. MXVTT is more reduced than MXBRGM with Fe(II)/Fe(III) ratio of 0.3 vs. 0.1. 2.1.2 Arrangements used in the sorption experiments Because iron was needed as Fe(II) in the sorption experiments, the iron in the solutions had to be reduced before use. The reduction was first tested with inactive FeCl3 solution having a similar NaCl concentration and pH as the radioactive solution in the planned sorption experiments. Reduction was carried out in an anaerobic glove-box by bubbling the solution with an H2/N2 gas mixture of 10/90 volume% using a Pt net as a catalyst. Figure 2 presents the increase of the Fe(II) concentration as a function of the bubbling time. As can be seen, Fe(III) is reduced in about seven hours to Fe(II). The radioactive isotope (55Fe) was used as a tracer in the sorption experiments. The half-life of Fe-55 is 2.7 years. The delivered radioactive Fe-55 was as FeCl3 dissolved in 0.1 M HCl. The total activity was 40 MBq (31.3.2009) and the activity concentration 630 MBq/cm3. The specific activity was 2 367 MBq/mgFe. The volume of the solution was 0.064 cm3 and its Fe concentration 0.2641 mg/cm3. The initial solution was first moved into 10 ml of 0.1 M HCl. This stock solution was used to prepare the reduced Fe(II) solutions, which were further used to prepare different types of sorption solutions. The method tested with the inactive solution was used to reduce the radioactive iron solutions. Typically, one millilitre of the stock solution was diluted to 100 ml of water or different solutions (depending on the planned experiment). Fe(III) or Fe(II) carrier

Reduction of Fe(III) in 0.3 M NaCl, 0.002 M HCl

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Figure 2. Increase of Fe(II) concentration as a function of bubbling time when Fe(III) was reduced by bubbling the solution with H2/N2 gas mixture of 10 /90 volume% using Pt net as a catalyst.

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was added to the solution so that the Fe concentration in the diluted solution was 10 µmol/l. This solution was then bubbled with H2/N2 or H2/Ar gas mixture using the platinum net as a catalyst. The tracer solutions for the sorption experiments were prepared from the reduced tracer solutions by diluting them to the ratio 1:10 with the background solution, 0.3 M NaCl, 0.05 M NaCl or acetate buffer of pH 5, depending on the planned conditions of the experiment. Strong NaOH, HCl or Fe(II) solutions were used to adjust the pH and Fe-background of the experiments when needed. The amounts of the solutions used for adjustment were so small that their effect on the total volume could be ignored. The clay contents in the sorption experiments were 1 g/l. For preparation of the sorption samples the bentonites dispersed to deionized water in concentration 10 g/l were mixed into the sorption solution to the ratio 1:10 in order to obtain the final clay concentration of 1 g/l. The sorption experiments were performed in centrifugal tubes of 30 ml. During the experiment, the sorption samples were closed in a steel vessel in the glove-box in order to minimize the risk of oxidation. In order to consume any oxygen in the metal vessel, in the first experiments it was filled with an H2/N2 (10/90 %) gas mixture together with a palladium catalyst. The vessel was shaken in the glove-box for about a week. In some cases parallel samples were kept in the glove-box atmosphere. No difference could be seen between the samples kept in the metal vessel and in the glove-box atmosphere, however. It seems that the metal vessel is not necessary but beneficial to ensure the condition during some accidental oxygen increase. In the later experiments, H2/N2 and palladium catalyst were no longer used in the metal vessel. The separation of bentonite and solution at the end of the sorption experiment was carried out in the earlier phase by centrifuging only. Since there was an indication that some colloidal bentonite may stay in the solution, the separation method was improved in the later phase of the study. The supernatants from the centrifuging were then ultrafiltered in order to remove the colloidal bentonite from the solution. It appeared that part of the earlier studies had to be omitted. The measurement of the Fe-55 was tested with a liquid scintillation analyzer. The samples were prepared from 3 ml of the radioactive solution and 17 ml of scintillation liquid (Perkin Elmer OptiPhase ‘HiSafe’ 3). The measurement efficiency of the arrangement was approximately 18 %. The background in the measurement was about 37 cpm. The Fe(II) and Fe(tot) analyses in the solutions were performed by spectrophotometry with the ferrozin method. In the case of Fe(II), the ferrozin buffer was added directly in the anaerobic glove-box to the samples. In the case of Fe(tot), thioglycolic acid was added into the sample in the glove-box. Then the sample was taken out, and heated to 90 oC for 30 minutes. Then the sample was cooled and the ferrozine buffer was added. The pH during the measurement was 4.1. When the concentration goes below 5 µg/l the result is rather scattered. Those values are presented only as <5µg/l and not used in the figures or calculations.

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2.1.3 Sorption experiments in 0.3 M NaCl at varying pH The sorption measurements were carried out in 0.3 M NaCl solution at varying pH. A suitable amount of the reduced Fe-55 tracer solution was further diluted so that Fe-55 activity was 3 – 4 kBq/ml, NaCl concentration 0.333 mol/l, Fe concentration 1 µmol/l and pH 3.7 – 3.9. The NaCl concentration of 0.333 M was used in order to compensate for the later dilution caused by adding the bentonite dispersion so that the final concentration was 0.3 M. The solution was once again bubbled with H2/N2 gas on a Pt-catalyst overnight to reduce possible trace amounts of Fe3+. Then the solution was used to prepare the sorption samples. In each sample tube was pipetted 13.5 ml of the tracer solution, 1.5 ml of bentonite dispersion in deionized water and different amounts of strong NaOH or HCl solutions to adjust the pH. Reference samples without bentonite were prepared by replacing the bentonite dispersion with 1.5 ml of water. After shaking for six days the sample tubes of the earlier experimental series were centrifuged, and the subsamples were taken from the supernatant for the pH measurement, Fe(II) and Fe(tot) analyses and Fe-55 activity measurements. Appendices A/1 and A/2 present the conditions and results of the earlier experiments. The results of the first experimental series indicated that colloidal bentonite may remain in the solution when centrifuging only was used for separation of the solution from the bentonite. New sorption experiments were carried out using an improved separation method. Principally, similar sorption samples as in the first series were prepared. After shaking for six to seven days the sample tubes were first centrifuged and a subsample was taken from the supernatant for the pH measurement. Then the rest of the supernatant was ultrafiltered with 10k MWCO filter. Fe(II) and Fe(tot) analyses and Fe-55 activity measurements were performed on the ultrafiltered solution. Appendices A/3 and A/4 present the conditions and results of the later experiments. 2.1.4 Sorption experiments in 0.05 M NaCl at varying pH The sorption measurements in 0.05 M NaCl were carried out in principle in the same way as those in the 0.3 M NaCl solution. A suitable amount of the reduced Fe-55 tracer was further diluted so that the Fe-55 activity was 3 – 4 kBq/ml, NaCl concentration 0.055 mol/l, Fe concentration 1 µmol/l and pH 3.7 – 3.9. The diluted solution was once again bubbled with H2/N2 gas on a Pt-catalyst overnight to reduce possible trace amounts of Fe3+. Then the solution was used to prepare the sorption samples. In each sample tube was pipetted 13.5 ml of the tracer solution, 1.5 ml of bentonite dispersion in deionized water and different amounts of strong NaOH or HCl solution to adjust the pH. Reference samples without bentonite were prepared by replacing the bentonite dispersion with 1.5 ml of water. After shaking for six days, the sample tubes of the earlier experiments were centrifuged and subsamples were taken from the supernatant for the pH measurement, Fe(II) and Fe(tot) analyses and Fe-55 activity measurements. Appendices B/1 and B/2 present the conditions and results of the earlier experiments. In the same way as in the 0.3 M NaCl solution the results of the first experimental series in 0.05 M NaCl solution indicated that there may stay colloidal bentonite in the solution

11

when centrifuging only was used for separation of the solution from the bentonite. New sorption experiments were carried out using an improved separation method. Principally similar sorption samples as in the first series were prepared. After shaking for six to seven days the sample tubes were first centrifuged and a subsample was taken from the supernatant for the pH measurement. Then the rest of the supernatant was ultrafiltered with 10k MWCO filter. Fe(II) and Fe(tot) analyses and Fe-55 activity measurements were performed on the ultrafiltered solution. Appendices B/3 and B/4 present the conditions and results of the later experiments. 2.1.5 Sorption experiments in 0.3 M NaCl solution at constant pH 5 in varying

Fe concentrations The experiments at a constant pH of 5 were carried out in 0.3 M NaCl solutions buffered with acetate buffer solution. The added concentration of the Fe(II) was varied in these experiments. A suitable amount of the reduced Fe-55 tracer solution was diluted so that the tracer activity was 3 – 4 kBq/ml, NaCl concentration 0.333 mol/l, acetate concentration 0,0108 mol/l and pH 5. The solution was once again bubbled with H2/N2 with a Pt net in order to ensure that all the iron was Fe(II). To adjust the iron concentration, Fe(II) carrier solutions of different concentrations were prepared in 0.333 M NaCl. In each sample tube was pipetted 13.5 ml of the tracer solution, 1.5 ml of bentonite dispersion in deionized water and 150 µl of Fe(II) carrier of different concentrations. Reference samples without bentonite were prepared by replacing the bentonite dispersion with 1.5 ml of water. After shaking for six days, the sample tubes were centrifuged. A subsample was taken from the supernatant for pH measurement. Then the rest of the supernatant was ultrafiltered, and subsamples were taken for Fe(II) and Fe(tot) analyses and Fe-55 activity measurements. Appendices C/1 and C/2 present the conditions and results of the experiments. 2.1.6 Results of the sorption experiments The sorption experiments were carried out with two batches of purified MX-80 bentonite. One was purified at BRGM France (MXBRGM) and the other one at VTT Finland (MXVTT). Experiments were also carried out with synthetic smectite, which did not include iron, prepared at BRGM (SBBRGM). Reference tests were carried out without any clay in the centrifugal tube (CT). The sorption experiments were carried out under low-oxygen conditions in an anaerobic glove-box in order to avoid oxidation of Fe(II) to Fe(III). Radioactive isotope (55Fe) was used as a tracer in the experiments. The experiments were carried out in 0.3 M and 0.05 M NaCl solutions as a function of pH, and in 0.3 M NaCl solution buffered at pH 5 as a function of the added Fe(II) concentration. The characterization of the clays shows certain differences between the properties of the different clays. The CEC value of MXVTT (0.975 meq/g) was higher than that of MXBRGM (0.791 meq/g) or SBBRGM (0.867 meq/g). MXVTT was also more reduced than MXBRGM with an Fe(II)/Fe(III) ratio of 0.3 vs. 0.1. In addition, the pH behaviour of the

12

clays varied somewhat, as seen in Figure 3 and Figure 4, which present the pH behaviour of the clays in the 0.3 M and 0.05 M NaCl, respectively. The figures present the pH in the solutions at the end of the test as a function of the NaOH amount used for the pH adjustment. Since the addition of bentonite dispersion (without NaOH) to the sorption solution already increases the pH, the curve without clay has been moved to start from the same pH as the curves with the clay. The pH of MXVTT increases more quickly with added NaOH than that of MXBRGM or SBBRGM indicating lower buffering capacity. The edge site capacities from pH 4.2 to 7 evaluated on the basis of the NaOH consumption difference between the curve without clay and with clay are 0.12–0.13 mmol/g for MXBRGM and SBBRGM and 0.08–0.09 mmol/g for MXVTT. Possible reasons for these differences could be, for example, the differences in the initial clay material used in the purification, or differences in the purification process. Behaviour of the inactive iron in the sorption experiments as a function of pH is presented in Figure 5 and Figure 6 for 0.3 M and 0.05 M NaCl solutions. The methods used for separation of the solid material, i.e. centrifugation (c) in the earlier experiments and centrifugation and ultrafiltration (c+u) in the later experiments, are given in the legends.

Experiment in 0.3 M NaCl

0

1

2

3

4

5

6

7

8

9

10

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

Added NaOH (µmol)

pH

MX(BRGM)

MX(VTT)

SB(BRGM)

No clay

Figure 3. pH behaviour of the sorption experiments in 0.3 M NaCl due to adding of NaOH. The sample volumes were 15 ml and their clay contents 1 g/l. For comparison the values for the solution without clay are given.

13

Experiment in 0.05 M NaCl

0

1

2

3

4

5

6

7

8

9

10

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

Added NaOH (µmol)

pH

MX(BRGM)

MX(VTT)

SB(BRGM)

No clay

Figure 4. pH behaviour of the sorption experiments in 0.05 M NaCl due to adding of NaOH. Dispersion volumes were 15 ml and their clay contents were 1 g/l. For comparison the values for the solution without clay are given.

The concentration of the added iron varied somewhat in the tests, being in the earlier tests 56 µg/l and in the later tests 61µg/l. The difference can be seen in the concentrations of the experiments without clay at low pHs, where no sorption occurs. No essential difference is seen between Fe(tot) and Fe(II) concentration in the experiments without clay, which indicates that iron is as Fe(II). At pH values above 6, the iron concentrations start to decrease, which indicates sorption on the tube walls or precipitation. In the experiments with MXBRGM and MXVTT the iron concentration in the solution increased strongly at low pH owing to the dissolution or desorption from the bentonite. The iron concentration does not seem to depend on the NaCl concentration used in the experiment. There is a clear difference between the Fe(tot) and Fe(II) concentrations caused by MXBRGM and MXVTT. MXBRGM causes a higher total iron concentration than MXVTT. In the tests with MXBRGM there is also a clear difference between the Fe(tot) and Fe(II) concentrations, which indicates that there is both Fe(III) and Fe(II) in the solution. In the tests with MXVTT the concentrations are lower and the Fe(tot) and Fe(II) concentrations are practically equal, which indicates that only Fe(II) is released into the solution. The iron concentrations of the MXBRGM tests decrease more quickly with increasing pH than those of the MXVTT tests such that already at pH 4 the concentration of the MXBRGM tests are smaller than the concentrations of the MXVTT tests. This is explained at least partly by the strong sorption of Fe(III) in the case of MXBRGM. The iron concentrations of the test series without ultrafiltering were higher than in the later series, where the improved separation was used. This suggests that part of the iron was bound in the bentonite colloids which were separated when the ultrafiltering was added in the separation. In the case of SBBRGM the iron concentrations in the solutions at low pH are close to the added concentration. No clear difference is seen between the two separation procedures. The concentration decreases owing to sorption when the pH increases over six.

14

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9 10

Final pH

Fe

tot

(µg

/l)

No clay (c)

No clay (c+u)

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9

Final pH

Fe(

II)

(µg

/l)

No clay (c)

No clay (c+u)

0

200

400

600

800

1000

2 3 4 5 6 7 8 9

Fe t

ot(µ

g/l)

Final pH

MX-BRGM (c)

MX-BRGM (c+u)

0

200

400

600

800

1000

2 3 4 5 6 7 8 9

Fe

(II)

(µ

g/l)

Final pH

MX-BRGM (c)

MX-BRGM (c+u)

0

100

200

300

400

500

600

2 3 4 5 6 7 8 9

Fe t

ot(µ

g/l)

Final pH

MX-VTT (c)

MX-VTT (c+u)

0

100

200

300

400

500

600

2 3 4 5 6 7 8 9

Fe(

II) (

µg

/l)

Final pH

MX-VTT (c)

MX-VTT (c+u)

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9

Fe t

ot(µ

g/l)

Final pH

SB-BRGM (c)

SB-BRGM (c+u)

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9

Fe(

II)

(µg

/l)

Final pH

SB-BRGM (c)

SB-BRGM (c+u)

Figure 5. Fe(tot) and Fe(II) concentrations in the solution as a function of pH at the end of the experiment with different clays (MXBRGM, MXVTT, SBBRGM) and without clay in 0.3 M NaCl. The separation methods, centrifuging (c) and ultrafiltering (u) are given in the legend.

15

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9

Final pH

Fe

tot

(µg

/l)

No clay (c)

No clay (c+u)

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9

Final pH

Fe

(II)

(µ

g/l)

No clay (c)

No clay (c+u)

0

100

200

300

400

500

600

700

800

2 3 4 5 6 7 8 9

Fe t

ot(µ

g/l

)

Final pH

MX-BRGM

MX-BRGM(c)

0

100

200

300

400

500

600

700

800

2 3 4 5 6 7 8 9

Fe(

II) (

µg

/l)

Final pH

MX-BRGM (c)

MX-BRGM (c+u)

0

100

200

300

400

500

600

700

800

2 3 4 5 6 7 8 9

Fe

tot(µ

g/l)

Final pH

MX-VTT (c)

MXVTT (c+u)

0

100

200

300

400

500

600

700

800

2 3 4 5 6 7 8 9

Fe

(II)

(µ

g/l

)

Final pH

MX-VTT (c)

MX-VTT (c+u)

0

10

20

30

40

50

60

70

2 3 4 5 6 7 8 9

Fe t

ot(µ

g/l)

Final pH

SB-BRGM (c)

SB-BRGM (c+u)

0

10

20

30

40

50

60

70

2 3 4 5 6 7 8 9

Fe(

II) (

µg

/l)

Final pH

SB_BRGM (c)

SB-BRGM (c+u)

Figure 6. Fe(tot) and Fe(II) concentrations in the solution as a function of pH at the end of the experiment with different clays (MXBRGM, MXVTT, SBBRGM) and without clay in 0.05 M NaCl. The separation methods, centrifuging (c) and ultrafiltering (u) are given in the legend.

16

Figure 7, Figure 8 and Figure 9 present the results for the radioactive Fe-55 tracer in the sorption experiments with 0.3 M NaCl as a function of pH. The distribution coefficient in the Figures was calculated from Eq. (1).

w

b

bd A

A

mA

VAAK

1

10

Equation 1

where Kd is the distribution coefficient, V is the volume of the solution in the sorption experiment, mb is the weight of bentonite, A0 and A1 are the total activities in the solution at the beginning and end of the test, and Ab and Aw are the tracer concentrations in the bentonite and in the solution at the end of the test. Qualitatively, the sorption of the radioactive Fe-55 on all the clays shows the same type of behaviour, i.e. sorption increases with increasing pH. In the measurements with the bentonite purified at VTT (MXVTT), the sorption occurs at a higher pH than in the measurements carried out with bentonite purified at BRGM (MXBRGM). In Figure 7, where the fraction of the sorbed Fe-55 is presented as a function of pH, the effect of the ultrafiltration is not seen very clearly, but in Figure 8 where the sorption has been presented as the distribution coefficient Kd, the effect can be seen. In the experiments with MXBRGM and MXVTT the Kd values increase at high pH when ultrafiltering is used, and the activity sorbed on bentonite colloids is removed from the solution. At pH values below 6, the Kd values with and without ultrafiltering are close to each other, however. In the case of SBBRGM the Kd values with and without ultrafiltering are rather similar at all pH values, indicating lower sorption where the small amount of bentonite left in the solution does not disturb. Figure 9 presents the distribution coefficient of Fe-55 for the experiments in 0.3 M NaCl as a function of pH when the results of the insufficient separation above pH 6 have been omitted. The omitted results are also indicated in Appendices A/1 – A/4. At low pH, the Kd values of MXVTT and SBBRGM are almost equal. Close to a pH of 6, the Kd of MXVTT increases quickly and starts to follow the Kd of MXBRGM. The differences between the results have also been obtained earlier by Tournassat (2003) when using bentonites of different batches. The differences were attributed to the reduction (iron/manganese oxide removal) and oxidation (organic matter removal) during the bentonite purification process. The results from the experiments in 0.05 M NaCl are presented in Figure 10, Figure 11 and Figure 12. The effect of the ultrafiltration is seen in the same way as in the experiments with 0.3 M NaCl. In the experiments with MXBRGM and MXVTT the Kd values increase at high pH when ultrafiltering is used and the activity sorbed on bentonite colloids is removed from the solution. At pH values below 6, the Kd values with and without ultrafiltering are close to each other, however. In the case of SBBRGM the Kd values are slightly higher with ultrafiltering at high pH, which may indicates the higher sorption and worse separation than in the experiments with 0.3 M NaCl. Figure 12 presents the distribution coefficient of Fe-55 for the experiments in 0.05 M NaCl as a function of pH when the results of the insufficient separation above pH 6 have been omitted. The omitted results are also indicated in Appendices B/1 – B/4. At low pH, the Kd values of MXVTT and SBBRGM are almost equal. Close to a pH of 6, the Kd of MXVTT increases quickly and starts to follow the Kd of MXBRGM.

17

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%)

MX-BRGM (c)

MX-BRGM (c+u)

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%)

MX-VTT (c)

MX-VTT (c+u)

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%) SB-BRGM (c)

SB-BRGM (c+u)

0

10

20

30

40

50

60

70

80

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%)

No clay (c)

No clay (c+u)

Figure 7. Sorbed fraction of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.3 M NaCl solution with different clays (MXBRGM, MXVTT, SBBRGM) and without clay. The separation methods, centrifuging (c) and ultrafiltering (u), are given in the legend.

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd (

ml/

g)

MX-BRGM (c)

MX-BRGM (c+u)

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd (

ml/

g)

MX-VTT (c)

MX-VTT (c+u)

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd (

ml/

g)

SB-BRGM (c)

SB-BRGM (c+u)

Figure 8. Kd of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.3 M NaCl solution with different clays (MXBRGM, MXVTT, SBBRGM) and without clay. The separation methods, centrifuging (c) and ultrafiltering (u), are given in the legend.

18

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd (

ml/

g)

MX-BRGM

MX-VTT

SB-BRGM

Figure 9. Kd of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.3 M NaCl solution with 1 g/l of different clays (MXBRGM, MXVTT, SBBRGM). The values of insufficient separation have been omitted (see Appendix A).

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%)

MX-BRGM (c)

MX-BRGM (c+u)

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%) MX-VTT (c)

MX-VTT (c+u)

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%) SB-BRGM (c)

SB-BRGM (c+u)

0

20

40

60

80

100

120

2 3 4 5 6 7 8 9 10

Final pH

So

rbed

Fe-

55 (

%) No clay (c)

No clay

Figure 10. Sorbed fraction of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.05 M NaCl solution with different clays (MXBRGM, MXVTT, SBBRGM) and without clay (CT). The separation methods, centrifuging (c) and ultrafiltering (u), are given in the legend.

19

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd

(m

l/g

)

MX-BRGM (c)

MX-BRGM (c+u)

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd

(m

l/g

)

MX-VTT (c)

MX-VTT (c+u)

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd

(m

l/g)

SB-BRGM (c)

SB-BRGM (c+u)

Figure 11. Kd of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.05 M NaCl solution with different clays (MXBRGM, MXVTT, SBBRGM) and without clay. The separation methods, centrifuging (c) and ultrafiltering (u), are given in the legend.

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

1.0E+07

1.0E+08

2 3 4 5 6 7 8 9 10

Final pH

Kd (

ml/

g)

MX-BRGM

MX-VTT

SB-BRGM

Figure 12. Kd values of the Fe-55 tracer as a function of the final pH in the experiments performed in 0.05 M NaCl solution with 1 g/l of different clays (MXBRGM, MXVTT, SBBRGM). The values of insufficient separation have been omitted (see Appendix B).

20

If the iron in the solution is in one redox state and the radiotracer behaves like the inactive iron, the estimation of Fe participating in the sorption/desorption process can be obtained from Eq. 1 by considering the concentration of natural Fe in solution together with the Kd value obtained for 55Fe. In the performed experiments the situation was more complex, however, when the iron could exist both as Fe(II) and Fe(III), and Fe-55 was added to the experiment as Fe(II). The kinetics of the isotope exchange between Fe(II) and Fe(III) is quick, however (Welch et al. 2003) and also the sorption/desorption can be assumed to occur quickly compared to the experimental time. It can thus be assumed that a complete isotope exchange occurs between Fe-55 and inactive Fe(II) and Fe(III), and after some time there are similar isotope distributions in the dissolved iron and sorbed iron of both redox states. In such a case, the average Kd obtained from the tracer measurement depends on the Fe(II) to Fe(III) ratio in the experiment and the separate Kd values of the different redox states. If the KD values of Fe(II) and Fe(III) are equal, it is possible to use the Kd of the radiotracer directly in the calculation of Fetot and different redox states participating in the sorption/desorption process. If the Kd of Fe(II) and Fe(III) are not equal, it is possible to calculate only the total concentration of iron participating in the sorption/desorption process from Fetot in the solution and average Kd given by the tracer.

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

0 1 2 3 4 5 6 7 8

Mo

bile

Fe t

ot(m

g/g

ben

t)

Final pH

0.3 M NaCl

MX-BRGM

MX-VTT

SB-BRGM

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

0 1 2 3 4 5 6 7 8

Mobile Fe(II) (mg/g

clay)

Final pH

0.3 M NaCl

MX‐BRGM

MX‐VTT

SB‐BRGM

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

0 1 2 3 4 5 6 7

Mo

bile

Fe t

ot(m

g/g

ben

t)

Final pH

0.05 M NaCl

MX-BRGM

MX-VTT

SB-BRGM

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

0 1 2 3 4 5 6 7 8

Mobile

 Fe/(II) (mg/g c

lay)

Final pH

0.05 M NaCl

MX‐BRGM

MX‐VTT

SB‐BRGM

Figure 13. Evaluation of the amount of the mobile inactive iron participating in the sorption-desorption process in the experiments with 0.3 M NaCl (upper) and 0.05 M NaCl (lower). The evaluation is based on the Kd, determined with the radioactive Fe-55 tracer, and on the Fetot and Fe(II) concentration in the solution.

21

The calculations of iron participating in the sorption process were done for the lower pH values where the iron concentration in the solution at the end of the sorption experiment was high enough to allow reliable analysis in the solution. Figure 13 presents the calculated amounts of Fetot and Fe(II) per the weight of bentonite as a function of the pH for the experiments in 0.3 and 0.05 M NaCl. The iron concentrations of MXBRGM and MXVTT are high at low pH, indicating dissolution of iron from bentonite. At higher pH the concentration decreases, approaching the value of the synthetic smectite which does not contain iron so that iron comes from the added iron in the sorption solution. The NaCl concentration of the solution has only a small effect on the result. In the case of MXBRGM there is both Fe(II) and Fe(III) in the solution and it can be assumed that Fe(III) sorbs more than Fe(II). Consequently the result is reliable for Fetot only in the case of MXBRGM. In the case of MXVTT the Fetot is practically equal with Fe(II) concentration in the solution and one can use average KD given by the Fe-55 experiment as a KD of Fe(II). Figure 14 - Figure 20 present the results for the sorption experiments performed in the acetate buffer of pH 5 in 0.3 M NaCl. The results are presented as a function of the added iron concentration for different clays and for the samples without clay. The sorption experiments with the radioactive Fe-55 tracer in the acetate buffer of pH 5 show decreasing sorption of Fe-55 as a function of the increasing concentration of the added Fe(II) (Figure 14). With MXBRGM the sorption percentage decreases from about 97 to 86 % and with MXVTT and SBBRGM from 30 – 40 % to about 15 % when the added Fe(II) increases from about 50 µg/l to 550 µg/l, respectively. The Kd values of MXBRGM vary from 3·104 to 3·103 and those of MXVTT and SBBRGM from 6·102 to 1.5·102 ml/g (Figure 15). In the case of MXBRGM the final Fe(II) and Fe(tot) values are clearly lower than the added Fe concentrations, which indicate sorption of the added inactive iron (Figure 16). In the case of MXVTT the iron concentrations in the solutions at the end of the experiments are higher than the added concentrations, which indicate the dissolution of iron from the bentonite and low sorption (Figure 17). With the synthetic smectite (SBBRGM), the final iron concentrations at low added Fe concentration are about equal with the added concentration (Figure 18). At higher values of the added iron, the measured concentrations are below the added concentration indicating weak sorption. In the tests without clay, the iron concentrations at the end are close to the added iron concentration with some outliners (Figure 19). Figure 20 presents the calculated amounts of Fe(II) and Fetot participating in the sorption process per the weight of bentonite as a function of the added iron for different clays. The iron concentration of the synthetic bentonite (SBBRGM) is equal with the added one. In the case of MXVTT the calculated Fe(II) and Fetot concentrations are higher than the added one, which indicates dissolution of iron from bentonite during the sorption experiment. In the case of MXBRGM the results are quite scattered.

22

-20

0

20

40

60

80

100

120

0 200 400 600

Added iron (µg/l)

So

rbe

d F

e-5

5 (

%)

MX-BRGM

MX-VTT

SB-BRGM

No bent.

Figure 14. Sorbed Fe-55 as a function of added iron concentration in the sorption solution in 0.3 M NaCl solution buffered at pH 5 with acetate buffer. The different clays (1 g/l) used in the experiment (MXBRGM, MXVTT, SBBRGM) and the experiment without clay are given in the legend.

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

0 100 200 300 400 500 600

Added Fe(II) (µg/l)

Kd o

f F

e-5

5 (

ml/g

)

MX-BRGM

MX-VTT

SB-BRGM

Figure 15. Kd of Fe-55 in the sorption experiment with 0.3 M NaCl at pH 5 as a function of the added Fe(II) concentration in the solution.

23

0

100

200

300

400

500

600

0 200 400 600

Added Iron (µg/l)

Fin

al i

ron

(u

g/l

) Fe(II)

Fe(tot)

Figure 16. Final iron concentration as a function of the added iron concentration in the sorption experiments with MXBRGM in 0.3 M NaCl at pH 5. The line presents a situation where the final iron is equal with the added one.

0

100

200

300

400

500

600

0 200 400 600

Added iron (µg/l)

Fin

al i

ron

(u

g/l

)

Fe(II)

Fe(tot)

Figure 17. Final iron concentration as a function of the added iron concentration in the sorption experiments with MXVTT in 0.3 M NaCl at pH 5. The line presents a situation where the final iron is equal with the added one.

24

0

100

200

300

400

500

600

0 100 200 300 400 500 600

Added iron (µg/l)

Fin

al i

ron

(u

g/l

) Fe(II)

Fe(tot)

Figure 18. Final iron concentration as a function of the added iron concentration in the sorption experiments with SBBRGM in 0.3 M NaCl at pH 5. The line presents a situation where the final iron is equal with the added

0

100

200

300

400

500

600

0 200 400 600

Added iron (µg/l)

Fin

al i

ron

(u

g/l

)

Fe(II)

Fe(tot)

Figure 19. Final iron concentration as a function of the added iron concentration in the sorption experiments without bentonite in 0.3 M NaCl at pH 5. The line presents a situation where the final iron is equal with the added one.

25

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 200 400 600

Mobile

 Fetot(m

g/g b

ent)

Added Fe(II) (µg/l)

MX‐BRGM

MX‐VTT

SB‐BRGM

0.0E+00

1.0E‐01

2.0E‐01

3.0E‐01

4.0E‐01

5.0E‐01

6.0E‐01

7.0E‐01

8.0E‐01

0 100 200 300 400 500 600

Mobile

 Fe(II) (mg/g b

ent)

Added Fe(II) (µg/l)

MX‐BRGM

MX‐VTT

SB‐BRGM

Figure 20. Evaluation of the amount of the mobile inactive Fetot (left) and Fe(II) (right) participating in the sorption-desorption process in the experiments with 0.3 M NaCl at pH 5. The evaluation is based on the Kd, determined with the radioactive Fe-55 tracer, and on the Fetot and Fe(II) concentration in the solution.

2.2 Diffusion experiments A non-steady-state diffusion experiment method, where the tracer is introduced as an impulse source between two bentonite plugs (Figure 21), was used in measurements of the diffusion of Fe(II) in bentonite. The cylindrical cell for the bentonite sample was made from PEEK. The diameter of the bentonite was 20 mm and the lengths of the two bentonite plugs, between which the radioactive tracer was placed, were 9 mm. The bentonite sample was closed between PEEK filter plates supported by perforated support plates made from PEEK. The diffusion sample was placed between two solution cells made from PVC. Membrane filters (mixed cellulose esters, pore size 0.45 µm) were placed between the bentonite and the PEEK filters to avoid release of colloidal bentonite to the solution. The samples were prepared and stored in the anaerobic glove-box. Parallel pH measurements were performed for following of pH in the bentonite during the experiments. Table 4 presents the experimental arrangements.

Figure 21. Diffusion cell for a non-steady-state diffusion experiment where the tracer is introduced as an impulse source between two bentonite plugs.

Solution cell

Membrane filter + PEEK filterplate

Perforated support

Tracer in membrane filter

Bentonite samples

26

Table 4. Experimental arrangements of the diffusion and pH experiments.

Sample code Concentration of NaCl in solution (mol/l)

Planned pH in bentonite/ solution

Planned experimental time (year)

Comments

Fe-8-1-D 0.3 8 1 Diff.sample Fe-8-2-D 0.3 8 2 - 4 Diff.sample Fe-8-1-pH 0.3 8 1 pH sample Fe-5-1-D 0.3 5 1 Diff.sample Fe-5-2-D 0.3 5 2 - 4 Diff.sample Fe-5-1-pH 0.3 5 1 pH sample 2.2.1 Preparation of the samples for the diffusion experiments at pH 8 Two diffusion samples were prepared for the experiments at a pH of about 8, one for the 1-year and the other one for the 2–4-year experiment. The samples were prepared from purified bentonite (MXVTT). For preparation of the two 9-mm-long plugs needed in the diffusion cell, two fully saturated 15-mm-long bentonite samples were first compacted from purified bentonite and deionized water. The bentonite plugs were then allowed to homogenize and equilibrate for 10 weeks in contact with the external solution in the same type of cells as seen in Figure 21. The vessels were filled with 30 ml of 0.3 M NaCl solution at pH 8-9 during the equilibration phase in order to get the pH adjusted close to 8. Then one of the pre-saturated samples was trimmed to a final length of 9 mm by removing extra bentonite from both ends. The sample was then introduced into one end of the 18 mm-long diffusion cell. A membrane filter was dipped in a tracer solution and placed on the end of the bentonite in the diffusion cell in order to introduce about 40 µl of the tracer solution homogenously over the surface of the bentonite. The pH of the tracer solution had been adjusted to 6, which was lower than the planned pH of 8 in the sample. The reason for using the lower pH was that there is a risk of sorption and precipitation increasing when the pH exceeds six, as was seen in the sorption experiments without bentonite. The bentonite was allowed to take care of the adjustment of the final pH in the bentonite. One end of the other pre-saturated bentonite plug was trimmed, and the bentonite was pushed into the diffusion cell so that the trimmed end was tight against the membrane. The extra bentonite from the outer end was cut away. The cell was closed between the end pieces and fixed between the solution cells which aimed to keep the samples well saturated and to maintain swelling pressure in the sample. The solution cells were filled with 30 ml of 0.3 M NaCl solution of pH 8. 2.2.2 Preparation of samples for the diffusion experiments at pH 5 Two diffusion samples were prepared also for the experiments at a pH of about 5. For the diffusion experiments at pH 5 the pH values of the bentonites had to be adjusted before compaction of the samples. For the pH adjustment, a buffer solution was prepared by dissolving 2 moles NaCl and 0.1 mole acetic acid into 900 ml of deionized water. The pH of the solution was then adjusted to 5 with about 83 ml of 1 M NaOH. The volume was finally adjusted to 1000 ml. Oxygen was removed from the buffer solution by argon bubbling and the solution was moved to an anaerobic glove-box.

27

Into six centrifugal tubes of 50 ml was weighed 9.7 g of purified bentonite (MXVTT), which had been in an anaerobic glove-box for two months in order to remove the oxygen. Into each centrifugal tube was added 30 ml of the acetate buffer. The bentonite and buffer solution were mixed carefully and allowed to react overnight. Then the centrifugal tubes were centrifuged and the supernatants (21-23 ml) replaced with new buffer solutions. Five treatments were needed until the pH in the removed solution was decreased to 5.1. Figure 22 presents the pHs at the end of each treatment. The bentonites from the six centrifugal tubes were moved to six squeezing cells where the solution was squeezed out until the wanted final density was obtained. The pHs of the squeezed solutions varied from 5.10 to 5.14. About two weeks was needed to squeeze the extra solution away. The 17.7-mm-long bentonite plugs from each squeezing cell were moved to the 15-mm-long equilibration tubes made from PEEK, and the excess of bentonite was cut away. The PEEK tubes were placed between the end pieces and fixed between the solution vessels. The vessels were filled with 30 ml of acetate buffer where the NaCl concentration was 0.3 M, acetate concentration 0.01 M and pH 5. The bentonites were allowed to equilibrate with the solutions so that the excess NaCl and acetate could leach out from the bentonites. After three weeks the solutions were changed and leaching was continued until the total time was 11 weeks. The equilibrated bentonite plugs were used for preparation of two diffusion samples. One of the equilibrated plugs was trimmed to the final length of 9 mm by removing extra bentonite from both ends. The sample was then introduced into one end of the 18 mm-long diffusion cell. A membrane filter was dipped in a tracer solution of pH 5 and placed on the end of the bentonite in the diffusion cell. One end of another equilibrated bentonite plug was trimmed, and the bentonite was pushed into the diffusion cell so that the trimmed end was tight against the membrane. The extra bentonite from the outer end was cut away. The cell was closed between the end pieces and fixed between two solution cells. The solution cells were filled with 30 ml of a solution of 0.3 M NaCl,

4

5

6

7

0 20 40 60 80 100 120 140

Cumulative volume of acetate buffer (ml)

pH

aft

er t

reat

men

t p

erio

d

Figure 22. The pH of the removed solutions at the end of repeated treatments of bentonites (9.7 g) with the acetate buffer of pH 5.

28

0.01 M acetate at pH 5. The radioactive isotope (55Fe) was used as a tracer in the sorption experiments. The half-life of Fe-55 is 2.7 years. The delivered radioactive Fe-55 was as FeCl3 dissolved in 0.1 M HCl. The total activity was 40 MBq (31.3.2009) and the activity concentration 630 MBq/cm3. The specific activity was 2 367 MBq/mgFe. The volume of the solution was 0.064 cm3 and its Fe concentration 0.2641 mg/cm3. The initial solution was first moved into 10 ml of 0.1 M HCl. This stock solution was used to prepare the reduced Fe(II) solutions, which were further used to prepare different types of sorption solutions. The method tested with the inactive solution was used to reduce the radioactive iron solutions. Typically, one millilitre of the stock solution was diluted to 100 ml of water or different solutions (depending on the planned experiment). Fe(III) or Fe(II) carrier 2.2.3 Storing and finishing of the diffusion experiments The diffusion samples were stored in an anaerobic glove-box (O2<1 ppm) where the cells were closed in metal vessels in order to avoid the effects of any accidental oxygen increase. After 310 days the short-term diffusion experiment of each pH was finished. The samples intended for the 2-4 year experiments were left to continue in the anaerobic glove-box. The bentonite of the finished experiment was pushed out step by step from the diffusion cell and slices were cut away until the membrane where the tracer was initially placed was reached. The weights of the slices were determined at once. The activity of each slice was dissolved in 2-3 ml portions of 0.5 M HCl in 0.3 M NaCl. The solutions of four subsequent leachings were mixed and evaporated to a final volume of 3 ml. The activity of the solution was measured with liquid scintillation method. In each experiment the activity was found only in two slices next to the membrane, i.e. the tracer was approximately within 0.5 mm from the initial source. The dry density of each sample based on the water content measurements is presented in Table 5. The apparent diffusivity was evaluated on the basis of the concentration profile in the sample from Eq. 2 (Yu and Neretnieks 1997).

tD

x

tDM

C

aa4

exp2

1 2

Equation 2

where C is the concentration at a distance x from the source layer, M is the total amount of tracer added to the source layer per unit area, and x is the distance from the source, t is time and Da is the apparent diffusion coefficient. By taking the logarithm of both sides of Equation 2, a linear expression is obtained as the function of x2, the slope of which gives the apparent diffusion coefficient, Eq. 3.

tD

x

tDM

C

aa42

1lnln

2

Equation 3

29

The apparent diffusivities were calculated from Eq. 3 using the data in Table 5. Due to the short diffusion distance only two points were available for the calculation of the slope. The thicknesses of the bentonite slices were evaluated on the basis of the weight, and the diffusion distance (x) was set at the midpoint of the slice. The apparent diffusivities are presented in Table 5 and are clarified in Figure 23 and Figure 24, which present the activity in bentonite slices vs. the distance from the impulse source. The measured activity distributions are compared with the profiles calculated from the Da given by the experiments (Table 5) and from Da values which are by a factor of ten larger or lower. The calculated curves using the measured Da describe the measured activity profiles rather well. Table 5. Results of the diffusion experiments. Sample Slice (mm) x (mm) Activity

(cpm/cm3) Da (m

2/s) Dry density (g/cm3)

Fe-5-1-D 0.0 – 0.17 0.085 327800

6.4·10-16 1.47 0.17 – 0.55 0.310 83730

Fe-8-1-D 0.0 - 0.127 0.064 567330

3.7·10-16 1.45 0.127 – 0.491 0.309 51500

0.0E+00

2.0E+05

4.0E+05

6.0E+05

8.0E+05

1.0E+06

1.2E+06

1.4E+06

0.0 0.2 0.4 0.6 0.8 1.0

Activity (cpm/cm

3)

Distance (mm)

Diffusion at pH 5

Da = 8.9 E‐17 m2/s

Da = 8.9 E‐16 m2/s

Da = 8.9 E‐15 m2/s

Measured

Figure 23. Measured activity profile in the diffusion sample at pH 5 and modelled profiles with different Da values.

30

0.0E+00

2.0E+05

4.0E+05

6.0E+05

8.0E+05

1.0E+06

1.2E+06

1.4E+06

1.6E+06

1.8E+06

0.0 0.2 0.4 0.6 0.8 1.0

Activ

ity (c

mp/

cm3)

Distance (mm)

Diffusion at pH 8

Da = 3.7 E-17 m2/sDa = 3.7 E-16 m2/sDa = 3.7 E-15 m2/sMeasured

Figure 24. Measured activity profile in the diffusion sample at pH 8 and modelled profiles with different Da values. 2.2.4 Comparison of measured and calculated diffusion coefficients The dependence between the apparent diffusivity (Da), effective diffusivity (De), pore diffusivity (Dp), porosity (ε), dry density (ρ) and sorption factor (Kd) can be presented as seen in Equations 4 and 5 (Yu and Neretnieks (1997).

Equation 4

Equation 5 Table 6 compares the measured apparent diffusivities with the ones calculated from the Kd of Fe-55 measured in this study. The Dp values needed in the calculation were taken from Sato (1992), who determined the Da (= Dp) values of tritiated water in compacted Kunipia bentonite as a function of dry density. The montmorillonite content of Kunipia was 95% which is close enough to our purified bentonite. The Dp of tritium was evaluated to be 1.2 ·10-10 m2/s at the density used in the diffusion experiments of this study. Both measured and calculated diffusivities show rather low values. No complete matching is obtained between the calculated and measured Da, however.

31

Table 6. Comparison of the calculated and measured Da values. Separation in the sorption measurement: c = centrifuge, u = ultrafiltration. pH of the diffusion sample

Kd (ml/g) Separation in sorption measurement

Calculated Da (m2/s)

Measured Da (m2/s)

pH 5.1 – 5.5 103 c or c+u 4.5·10-14 6.4·10-16

pH 8 – 7.5 105 c 5·10-16 3.7·10-16

pH 8 – 7.5 107 c+u 5·10-18 3.7·10-16 2.2.5 Measurement of pH in bentonite samples Parallel samples to the diffusion experiments were prepared for measurement of pH in bentonites during the short term diffusion experiment. The samples were prepared in the same way as the diffusion samples, but the membrane filter where the radioactive tracer was added did not contain any Fe-55 activity. Otherwise the samples were like those prepared for the diffusion experiments at pH 8 and 5. Two IrOx electrodes (Muurinen and Carlsson 2010, Yao et al. 2001) calibrated at pHs of 4, 7, 10 and 12 beforehand were placed in bentonites at a distance of 6 mm from each end. A hole was made in the bentonite to introduce the electrode. After inserting the electrodes, the bentonite was compressed gently in order to obtain a good contact between the electrode and bentonite. The solution cells were filled with similar solutions as used in the diffusion experiments. The reference electrodes were placed in the solution cells during the measurement periods. For determination of pH the potential between the IrOx electrode and the reference electrode was measured. The pH values were calculated from the calibration equations for each electrode. At the end of the experiment the IrOx electrodes were removed by washing the bentonite away with 1M NaCl solution and the electrodes were recalibrated. The calibration curves before and after the measurements are presented in Figure 25 and Figure 26.

32

y = -58.5x + 927

y = -56.6x + 885

0

100

200

300

400

500

600

700

800

0 5 10 15

Po

ten

tial

vs.

SH

E (

mV

)

pH

Ir2

9.5.2011

17.4.2012

y = -58.5x + 938

y = -57.7x + 904

0

100

200

300

400

500

600

700

800

0 5 10 15

Po

ten

tial

vs.

SH

E (

mV

)

pH

Ir3

9.5.2011

17.4.2012

Figure 25. Calibration of the pH electrodes Ir2 and Ir3 before and after the experiment. The electrodes were used in the sample of pH 5.

33

y = -58.5x + 888

y = -56.3x + 842

0

100

200

300

400

500

600

700

800

0 5 10 15

Po

ten

tial

vs

SH

E (

mV

)

pH

Ir4

9.5.2011

17.4.2012

y = -58.5x + 907

y = -56.4x + 857

0

100

200

300

400

500

600

700

800

0 5 10 15

Po

ten

tial

vs.

SH

E

pH

Ir5

9.5.2011

17.4.2012

Figure 26. Calibration of the pH electrodes Ir4 and Ir5 before and after the experiment. The electrodes were used in the sample of pH 8. The electrodes retained the slope rather well but the constant of the regression equation decreased. For calculation of the pH values as a function of time, a time-dependent calibration equation was used assuming that both the slope and constant of the calibration curve change linearly as a function of time. Figure 27 presents the pH values as a function of time based on the time-dependent calibration equation. It is obvious that the pH changes slowly during the experiment so that in the sample where the initial pH was about 8, the pH at the end of the experiment is about 7.3. In the sample where the initial pH was about 5.1 the final pH increased to 5.5.

34

4

5

6

7

8

9

10

0 100 200 300 400

pH

Time (day)

Ir2 Ir3

Ir4 Ir5

Figure 27. Measured pH as a function of the experimental time in the samples of pH 5 (electrodes Ir2 and Ir3) and pH 8 (electrodes Ir4 and Ir5).

35

3 MODELLING OF THE EXPERIMENTS 3.1 Data selection and corrections 3.1.1 Problem of solid / liquid separation Inactive iron is released from bentonite into the solution in the sorption experiments. At first sight, these results could be linked considering that the concentrations of natural Fe originate principally from Fe sorption sites. The higher the Fe affinity is for the surface, the lower the Fe concentration in solution at a given pH. With this hypothesis, a certain amount of Fe(II) and Fe(III) is bounded to the surface sites prior to the experiment in the same way that Bradbury and Baeyens (1997) considered the competition effect of naturally occurring Zn in their montmorillonite material. In addition to the desorption, the iron can originate also from the dissolution of the clay network and/or the dissolution of accessory minerals. Mössbauer spectra, shown in Figure 1, evidence indeed the presence of small amount of iron oxides in both VTT and BRGM MX80 samples. Some discrepancies exist in the Fe behaviour of MXBRGM and MXVTT. Solubilisation of natural Fe from the montmorillonite material is higher in the pH range 2 – 4 but lower in the pH range 4 – 6 for MXBRGM than for MXVTT. MXBRGM exhibits higher Fetot than Fe(II) concentration in the sorption solutions at low pH and this effect is more pronounced when centrifuging only is used for separation. Both MXBRGM and MXVTT show very scattered KD values at pH > 6 when centrifuging only is used for separation of bentonite. Figure 28 presents the activity of the Fe-55 in the solutions of the sorption experiments in 0.3 and 0.05 M NaCl as a function of pH. The separation of the bentonite with the centrifuge (c) and centrifuge plus ultrafiltering (c+u) are given in the legend. The behaviours are rather similar in all experiments. At low pH, the tests of the two separation methods show rather similar behaviour. When centrifuging only is used the activities stay about constant when the pH exceeds six, i.e. when about 97 % of the activity has been sorbed. In the case where centrifuging and ultrafiltering are used the activity continues decreasing at a pH above six, however. It is most probable that the difference is caused by the particles in the solution when the centrifuge only is used. Figure 29 presents theoretically how the activity behaves if different percentages of the solid material stay in the solution after separation. As can be seen, the behaviour in Figure 29 is very similar to that in Figure 28. Comparison of the measured values with the theoretical curves suggests that typically 1-2 % of bentonite stays in the solution after centrifugation. A separation by centrifugation only leads to the presence of remaining clay particles in the analyzed supernatant and a wrong estimate of solute Fe in solution but only a weak underestimation of the Kd if less than about 97 % has been sorbed. The values where the particles obviously disturb have to be omitted, however. At a high pH, only the results where ultrafiltering was used can be accepted. In most cases when the pH exceeds 6 the inactive iron is below the determination level and only the Kd of the radioactive experiments can be used. Moreover, discrepancies between solute Fe(II) and total solute

36

Fe are measured in solutions corresponding to “centrifugation” condition with MXBRGM material. Usually, these differences are interpreted as a measurement of the presence of Fe(III) in solution. However, these concentrations can be as high as 0.2 µmol/L at pH ~7, that is, far higher than the solubility limit of iron hydroxides (ferrihydrite and goethite, for instance). Even at pH 4.7, the solution would be supersaturated with regards to ferrihydrite. As a consequence, it is very unlikely that the measured difference between solute total Fe and solute Fe(II) corresponds to solute Fe(III). Given the presence of remaining particles in suspension after centrifugation, this difference must rather be attributed to the (partial) reductive dissolution of these particles containing Fe in their structure during solute Fetot measurements.

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

0 5 10

Fe‐55 in

 solution (cpm)

pH

MX‐BRGM 0.3 M NaCl

c

c+u

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

0 5 10

Fe‐55 in

 solution (cpm)

pH

MX‐VTT 0.3 M NaCl

c

c+u

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

0 2 4 6 8 10

Fe‐55 in

 solution (cpm)

pH

MX‐BRGM 0.05 M NaCl

c

c+u

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

0 2 4 6 8 10

Fe‐55 in solution (cpm)

pH

MX‐VTT 0.05 M NaCl

c

c+u

Figure 28. The analyzed concentrations of the Fe-55 tracer in the solution after separation of the bentonite with centrifuge (c) and centrifuge + ultrafiltering (c+u). Experiments with MXBRGM and MXVTT in 0.3 and 0.05 M NaCl.

37

1.0E+00

1.0E+01

1.0E+02

1.0E+03

1.0E+04

1.0E+05

1.0E+06

2.0E+00 2.0E+02 2.0E+04

Mea

sure

d a

ctiv

ity

(cp

m)

Real Kd (ml/g)

0.0 %

0.5 %

1.0 %

1.5 %

2.0 %

Figure 29. Calculated activity as a function of Kd in the sorption experiment when different percentages of the solid material are not separated from the solution. The solution-to-solid ratio and the initial activity are as used in the present experiments.

Consequently, we must consider for modelling purposes:

Measured Fe(II) concentrations are the most reliable solute Fe concentration measurements;

Above pH 5-6 only the Kd of the radioactive experiments can be used. 3.1.2 Desorption vs. dissolution at low pH In the case of negligible dissolution, mobile Fe (sum of sorbed + solute iron) must be a constant value. This is true for all pH values in the experiment with synthetic clay, since this clay did not contain any Fe in its structure (Figure 13). For MX80 (BRGM or VTT), mobile Fe concentration values at a pH above 4.5-5 should also be constant because clay dissolution is negligible in these conditions. Moreover, in these conditions, mobile Fe concentrations at 0.05 and 0.3 mol/L NaCl should also be identical as the starting clay material was the same. At a pH above 5, mobile Fe concentration values are largely scattered. This is due to the low precision of the Fe concentration measurement at very low Fe(II) concentrations (<1 µmol/L) by spectrophotometry in combination with the large values of Kd at pH >5. It is possible to obtain a constant mobile Fe(II) concentration for pH>4.5 by slightly adjusting the solute Fe(II) concentrations (Figure 30). Doing so, it is also possible to compare the “natural Fe Kd values” to the 55Fe Kd values. “Natural Fe Kd values” can be calculated from the difference between Fe concentration at very low pH and Fe concentration at the pH of interest:

38

solution

solution2~pHmob

bnat [Fe]

[Fe] - ][Fe

m

V Kd Equation 6

2 3 4 5 6 70

10-5

2x10-5

3x10-5

4x10-5

pH

Mo

bile

Fe(

II) c

on

cen

trat

ion

(m

ol L

-1)

MeasuredAdjusted

2 3 4 5 6 70

2x10-6

4x10-6

6x10-6

8x10-6

10-5

1.2x10-5

1.4x10-5

pH

So

lute

Fe(

II) c

on

cen

trat

ion

(m

ol L

-1)

MeasuredAdjusted

4.5 5 5.5 6 6.5 70

5x10 -7

10 -6

1.5x10 -6

2x10 -6

sorption/desorption

Scatte

r due

to p

recis

ion of

solut

e

Fe(II)

conc

. mea

s. an

d high K

d value

s

Dissolution

Figure 30. Correction of solute Fe(II) concentrations for experiments conducted at pH > 4.5. In the case of negligible dissolution, i.e. almost all solute iron originates from desorption, then the total concentration of labile (mobile = solution + surface) Fe (Femob) can be obtained from experiments at low pH where sorption is negligible ([Femob] = [Fe(II)]solution in these conditions). This value must be in turn commensurable with [Femob] estimated from [Fe(II)]solution and radiotracer Kd value for different experimental conditions (pH, ionic strength). On the contrary, dissolution process produces Fe(II) in solution that is in excess as compared to isotopically exchangeable surface Fe (note that for simplification reasons, we did not consider the presence of Fe(III) in the system; its consideration would make the calculation more difficult and the comparison between MXVTT, without Fe(III) in solution and MXBRGM, with Fe(III) in solution). If dissolution dominates over desorption, it should lead to a difference in radiotracer Rd and estimated natural Fe Rd. This effect can be seen in Figure 31.

39

1 2 3 4 5 6 7

0

1

2

3

4

5

6

3 4 5

0

1

2

3

1 2 3 4 5 6 7

0

5

10

15

20

25

3 4 5

0

1

2

3

Desorption

Dissolution

Total Fe in solution

Fe

conc

entr

atio

n (

u.a

.)

pH

Desorption = 1/5 Dissolution

Radiotracer

Natural Fe

log

Rd

(u.

a.)

pH

Desorption

Dissolution

Total Fe in solution

Fe

con

cen

tra

tion

(u.a

.)

pH

Desorption = 5 Dissolution

Radiotracer

Natural Fe

log

Rd

(u.

a.)

pH Figure 31. Illustration of the effect of the relative contributions of desorption and dissolution on the total Fe concentration on the distribution coefficient for natural Fe and Fe radiotracer. Figure 32 (bottom) shows that dissolution is almost negligible for MXBRGM when compared to desorption, whereas it is the contrary for MXVTT. For both clays, Fe dissolution at pH 2.13 amounts to ~6.5 10-6 mol/L (calculated from the difference between total mobile Fe concentration and initially sorbed iron plus added iron), showing that the difference is due to the mass of initially sorbed Fe and not in a difference of dissolution extent. In turn, radiotracer Rd and “natural Rd” values are different for the VTT sample while they are very similar for the BRGM sample (Figure 32). It must be noted that this “natural Fe Kd” evaluation method for Fe sorption is identical to that used in Tournassat (2003) when only data on natural iron was available. Figure 33 shows that results obtained in the present study with MXBRGM are in good agreement with the results obtained by Tournassat (2003) for similar [Femob].

40

Figure 32. Comparison of Kd values calculated from difference in natural Fe concentrations (“natural Fe Kd”) and from radiotracer analyses. Data obtained in NaCl 0.3 mol L-1 (upper figures) and 0.05 mol L-1 (bottom figure). For MXBRGM (right) natural Fe Kd was calculated based on Fe(II) solution concentration. Data from experiments A1 and B1.

41

Figure 33. Comparison of Kd obtained in the present study with MXBRGM (circles) and Kd values calculated from data of Tournassat 2003 obtained at high NaCl concentration (>0.5 mol/L) and low total Femob concentration. The colour scale is indicative of the total mobile iron concentration (sorbed + solution).

3.2 General considerations for modelling We face two distinct Fe sorption behaviours for the two investigated clay samples. In the first case (BRGM sample and data from Tournassat, 2003), Fe(II) originating from the clay material behaves like the radiotracer. In the second case (VTT sample), natural Fe(II) originating from the clay material exhibits an apparent higher affinity for the surface than the radiotracer. We demonstrated how important the relative contributions of initially sorbed Fe are over Fe dissolution from the sample. In the following sorption edge modelling work, the above determined initial sorbed amount is used in the mobile Fe inventory (7.4 10-6 mol/g for MXBRGM, 1.9 10-6 mol/g for MXVTT

and 0 mol/g for synthetic smectite).

42

3.3 Modelling Fe sorption on synthetic smectite with a cation exchange / surface complexation model

The problems explicated in the above section are not present in experiments with synthetic smectites. Fe(II) sorption models were thus tested first on this clay material. As a preliminary predictive approach, we applied the parameters given in Table 1 that were obtained with the LFER from Bradbury and Baeyens (2005) for Fe(II). The results are shown in Figure 34. Fe(II) sorption is clearly underestimated by this preliminary modelling approach.

Figure 34. Comparison of LFER predictive model (black circles) with Fe(II) sorption data (open circles) on synthetic smectite. Left: sorption edge at 0.3 mol L-1 NaCl. Middle: sorption edge at 0.05 mol L-1 NaCl. Right: sorption isotherm at 0.3 mol L-1 NaCl and pH 5. Figure 35 show that data obtained in the present study at pH > 6 can be correctly reproduced by increasing the Fe2+ affinity constant from -0.45 to 0.5 and increasing the FeOH+ affinity constant from -10.4 to -7. Still, sorption data below pH 5 (arrows in the figure) are not well predicted with difference in predicted vs. measured Rd of ~1 log unit. This difference is significant since it represents a sorption extent difference of ~25%. At first sight, this difference might be attributed to an inaccurate value for the Fe2+ exchange selectivity coefficient. The tabulated value (log Kexch = 0.27) lies in the low range of values available from the literature (Tournassat et al. 2009). Moreover, FeCl+ exchange was not considered in these simulations. We tested this hypothesis by raising the Fe2+ selectivity coefficient to 0.5 (in log scale) and by introducing an Na-FeCl+ exchange reaction with log Kexch = 2.1, these values being in agreement with the recommendations of Tournassat et al. (2009). Figure 36 shows that sorption edge at 0.05 mol L-1 NaCl is now correctly reproduced but sorption edge at 0.3 mol L-1 NaCl cannot be explained only by cation exchange reaction (the log of Fe2+ selectivity coefficient should have been raised to a value of 1.5 that (i) is not compatible with the data obtained at 0.05 mol L-1 NaCl and (ii) has never been reported in the literature.

43

Figure 35. Comparison of Fe(II) sorption data on synthetic smectite (open circles) with results from the LFER model with adjusted Fe(II) affinities for strong sites (black circles). Left: sorption edge at 0.3 mol L-1 NaCl. Middle: sorption edge at 0.05 mol L-1 NaCl. Right: sorption isotherm at 0.3 mol L-1 NaCl and pH 5. The arrows indicate zones of significant discrepancy between model and data.

Figure 36. Comparison of Fe(II) sorption data on synthetic smectite (open circles) with results from LFER model with adjusted Fe(II) affinities for strong sites and increased Fe (Fe2+ and FeCl+) exchange selectivity coefficient (black circles). Left: sorption edge at 0.3 mol L-1 NaCl. Middle: sorption edge at 0.05 mol L-1 NaCl. Right: sorption isotherm at 0.3 mol L-1 NaCl and pH 5. The arrows indicate zones of significant discrepancy between model and data. Clearly, an additional sorption mechanism is needed to explain the results obtained at pH < 5. This might be linked to the presence of an additional adsorption site with an even higher affinity for Fe(II) than the strong site considered in the above modelling work. This might also be due to an incorporation of Fe in the clay structure (e.g. substitution with octahedral Mg). In the following, the first hypothesis (sorption site) is explored. In a preliminary approach, we tried to find a minimal set of model parameters that enables reproducing Fe sorption curves on the synthetic smectite (“minimal model”).

44

We considered Bradbury and Baeyens’ model with its protonation/deprotonation constants for strong and weak sites except that we did not consider the protonation reaction of strong sites (Tournassat et al., in prep, showed that this reaction was not constrained at all and can be removed to simplify the model). Site densities were considered to be adjustable parameters. This hypothesis is justified owing to the nature of the investigated material, a synthetic smectite, whose reactive surface area might be different from those described in the paper by Bradbury and Baeyens. We found that a model with a cation-exchange site and two strong sites is sufficient to fit satisfactorily the whole set of data. Model parameters are given in Table 7, and model results are shown in Figure 37. The total surface density of a strong site is commensurable with the strong site density given in Bradbury and Baeyens’ model (2005), i.e. 4.3 vs. 2 mmol of site per kg. According to these authors, this strong site density is a rather ill-defined parameter (Bradbury and Baeyens, 1997). The affinity of Fe2+ for the strong site having the highest surface density (site ≡Sb) is in good agreement with its estimate made with the LFER (~-0.4 in log scale). However, the affinity of FeOH+ is significantly higher (-7.5 vs. -10.4) and is thus in disagreement with the LFER estimate. Concerning the site ≡Sa, its affinity for Fe2+ is clearly in disagreement with the LFER. It must be noted that the sorption edge at 0.3 mol L-1 NaCl can be reproduced with a very restricted range of values for the site ≡Sa density and its affinity for Fe2+.

Table 7. Parameters of “minimal model”. Protonation/deprotonation constants were taken from Bradbury and Baeyens (1997). Cation-exchange parameters are taken from the model shown on Figure 36.

Strong sites densities mol kg-1 ≡SaOH 3 10-4 ≡SbOH 4 10-3 Strong sites surface complexation for iron log K ≡SaOH + Fe2+ ↔ ≡SaOFe+ + H+ 3.5 ≡SbOH + Fe2+ ↔ ≡SaOFe+ + H+ -0.4 ≡SbOH + Fe2+ + H2O ↔ ≡SbOFeOH + 2H+ -7.5

Figure 37. Comparison of Fe sorption data on synthetic smectite with the “minimal model” results.

45

Model strong disagreements with the LFER predictions might be attributed to the Fe(II) surface oxidation process described by Gehin et al. (2007) on the same type of synthetic smectite. Gehin et al. (2007) evidenced a partial oxidation of Fe(II) on the surface of synthetic clay despite the absence of oxidant in the system other than water. This result was confirmed in Charlet et al. (2007) who showed that Se(IV) could be reduced by the clay with sorbed Fe, that the extent of Se(IV) reduction was commensurable with the amount of oxidized Fe (taking into account the Se(IV)Se(0) vs. Fe(II)Fe(III) electron stoichiometries), and that this Se(IV) reduction was shifted in time with regard to Fe(II) oxidation highlighting the need to consider an electron “storage” between the two redox processes. We attempted to model the Fe(II) sorption data of this report with the sorption/oxidation model of Gehin et al. (2007) but this model led to a predicted sorption extent far too high compared to the measured one in the present study (not shown). It was thus necessary to correct Gehin's model thanks to the new data available in the present report. 3.4 Modelling Fe sorption on synthetic smectite with a cation exchange /

surface complexation model, taking into account surface redox reactions

Data from Gehin et al. (2007) were reprocessed to constrain the proposed model. In the following, the models were run with the same parameters for the synthetic smectite used in the present study and for those used by Gehin et al. excepting the total cation-exchange capacity (CEC) : Gehin et al. measured a CEC of ~0.5 mol kg-1 instead of 0.87 mol kg-1 in the present study. Gehin et al. measured only one Fe2+ sorption curve as a function of pH in a CaCl2 ionic background. We considered that Ca and Fe have the same affinity for cation-exchange sites (Tournassat et al. 2009). Since Gehin et al.’s measurements were performed at high Fe surface loading (Fe/clay ratio amounted to 63 mmol kg-1), their Fe sorption edge was mostly representative of weak site reactivity, which is not probed by the new results presented in this report. However, the information gained from 57Fe Mössbauer spectrometry as a function of pH is mostly linked to strong site reactivity for pH < 7. These data can thus be used to constrain the proposed model. Gehin et al.’s model hypothesis is that Fe2+ surface oxidation in Fe3+ is driven by a mechanism of sorbed-Fe redox stabilization. In the following, we construct such a model, step by step, which can take into account information from the present report (sorption edges and isotherms at low total Fe concentration) and from the publication of Gehin et al. (sorption edge and surface Fe redox speciation at high total Fe concentration). While sorption edges from the present study evidence a strong Fe sorption increase as a function of pH in the pH range 2 to 6, data from Gehin et al. exhibit a constant sorption extent at a value corresponding to ~15 % of total iron. This amount of sorbed Fe cannot be due to cation exchange only, as shown by their Mössbauer results showing quadrupolar components that have different properties than cation-exchange species. This contrasting result between the present study and Gehin et al.’s study can be interpreted by a site saturation effect due to the high Fe concentration used in Gehin et al. The total site density can be estimated from Mössbauer results: it represents

46

approximately half of the total sorbed iron between pH 4 and 6, i.e. ~ 5 mmol kg-1. This value is in remarkable agreement with the total strong site density obtained in Table 7. Site saturation takes place in the whole investigated pH range. This feature is hardly rendered by a classical surface complexation reaction that depends on pH (see e.g. Table 7). The simplest surface reaction stoichiometry that allows such a behaviour is the following:

>SOH + Fe2+ ⇌>SOHFe2+ log K1 Reaction 1

This reaction can be seen either as a surface complexation reaction or as a substitution/growing mechanism on the clay surface. From pH 7, the percentage of sorbed Fe starts to increase in the experiment of Gehin et al. to attain almost 100 % at pH 8.5. This result can be modelled by considering the sorption of Fe2+ on weak sites with the stoichiometry:

>WOH + Fe2+ + 2 H2O ⇌>WOFe(OH)2- + 3 H+ log K2 Reaction 2

Figure 38 shows that sorption data from Gehin et al. (2007) are well reproduced by this very simple model with log K1 = 4.5 and log K2 = -19 and a weak site density of 0.08 mol kg-1 (hereafter named model A). The log K2 value (-19) is in remarkable agreement with the value estimated from the LFER (-19.8, see Table 1) although other species predicted by the LFER are unnecessary to reproduce the sorption curve. This model predicts the presence of Fe(II) species at the surface while Mössbauer measurements made by the authors evidenced the presence of Fe(III). Moreover, if we try to apply model A to the data obtained in the present study, sorption extents are strongly underestimated (not shown). It was thus decided to test if the introduction of surface oxidation processes enabled the reproduction of the results from both Gehin et al. and the present study.

2 3 4 5 6 7 8 90

20

40

60

80

100

pH

% s

orb

ed

Fe

Figure 38. Results from Gehin et al. (2007) (circles) with fitted model line according to hypotheses from model A.

47

The following reactions were introduced in the model for strong sites :

>SOH + Fe3+ ⇌>SOFe2+ + H+ log K3 Reaction 3

>SOH + Fe3+ + H2O ⇌>SOFeOH+ + 2 H+ log K4 Reaction 4

>SOH + Fe3+ + 3 H2O ⇌>SOFe(OH)3- + 4 H+ log K5 Reaction 5

Figure 39 shows that this simple model reproduces remarkably well the sorption data obtained in the present study at low Fe concentration, as well as Mössbauer speciation results obtained at strong site saturation conditions, giving confidence in the mechanism underlying the model: the very strong preference of the synthetic clay structure/surface for Fe(III) compared to Fe(II) drives the Fe(II) surface oxidation mechanism that is itself responsible for the clay strong sorption affinity for Fe(II). All of the parameters of the model (hereafter named Reference Synthetic Montmorillonite Model – RSM model) are given in Table 8. Although other models could certainly be proposed, it is worth noticing that the proposed model enables reproduction of the peculiar sorption behaviour of Fe with only one type of strong site, whereas it was necessary to introduce a second site with the “minimal model” that did not take into account the possible surface oxidation of Fe.

48

Figure 39. Results of RSM model (Table 8) compared to data from the present study and from Gehin et al. (2007).

49

Table 8. Parameters of reference synthetic montmorillonite model (RSM; non electrostatic model).

Cation Exchange site density (X-) 0.87 mol kg-1 Cation Exchange reactions log K 2 XNa + Fe2+ ↔ X2Fe + 2 Na+ 0.5 XNa + FeCl+ ↔ XFeCl + Na+ 2 2 XNa + Ca2+ ↔ X2Ca + 2 Na+ 0.5 XNa + CaCl+ ↔ XCaCl + Na+ 2 Strong sites density 5 10-3 mol kg-1 Strong sites surface reactions log K ≡SOH ↔ ≡SO- + H+ -7.9 ≡SOH + Fe2+ ↔ ≡SOHFe2+ 4.5 ≡SOH + Fe3+ ↔ ≡SOFe2+ + H+ 16 ≡SOH + Fe3+ + H2O ↔ ≡SOFeOH+ + 2 H+ 11 ≡SOH + Fe3+ + 3 H2O ↔ ≡SOFe(OH)3

- + 4H+ -4 Weak sites density 8 10-2 mol kg-1 Weak sites surface reactions log K ≡WOH ↔ ≡WO- + H+ -7.9 ≡WOH + H+ ↔ ≡WOH2

+ 4.5 ≡WOH + Fe2+ + 2 H2O ↔ ≡WOFe(OH)2

- + 3 H+ -19 3.5 Modelling structural iron redox properties in natural clays In a previous report, we have shown that Fe(II) sorption results from Tournassat (2003) could be simulated adequately with a model considering a two-steps uptake process with first Fe(II) adsorption on clay edges followed by surface oxidation by structural Fe(III). Recent chemical and spectroscopic investigations of Fe(II) sorption on clay minerals gave rise to a surface redox model (Hofstetter et al. 2006; Merola et al. 2006; Lerf et al. 2011; Schaefer et al., 2011). Using Mössbauer spectrometry coupled to wet chemical experiments with pure 56Fe (transparent to Mössbauer effect) and 57Fe solutions, Schaefer et al. have shown that sorbed Fe(II) was (i) oxidized, (ii) precipitated as lepidocrocite (their total concentration of Fe was high enough) and (iii) was responsible at the same time for structural clay Fe(III) reduction into Fe(II). Although their results were obtained on a ferruginous smectite (nontronite) and not on montmorillonite, their data can be used as a proxy in order to estimate thermodynamic parameters for the Fe surface/structural redox reactions. The hypothesized mechanism of Fe(II) adsorption and oxidation reaction is depicted in Figure 40.

50

Figure 40. Hypothesized mechanism for Fe(II) adsorption and surface oxidation by natural smectites. 1: specific sorption of Fe(II) (pH dependant). 2: electron transfer from sorbed Fe(II) to structural reactive Fe(III). 3: Specific desorption of Fe(III) (pH dependant). 4: iron hydroxide precipitation. The experimental observations made by Schaefer et al. (2011), when combined with the observations made by Gehin et al. (2007) give rise to a coherent picture of the processes acting during Fe(II) sorption on a natural montmorillonite. However, the redox properties of structural iron in clay are largely unknown. In the present report, we combined the RSM model proposed above to the result of Schaefer et al. in order (i) to reproduce Fe(II) sorption data and observations made by Schaefer et al. and (ii) to estimate a redox potential for clay structural iron. Schaefer et al. provided: (1) 56Fe sorption data on a nontronite smectite at high Fe/clay ratio (Figure 41). These data were digitized (Table 9); (2) a linear relationship between structural Fe(II) probed by Mössbauer spectrometry and Fe(II) uptake from the solution; (3) a quantitative speciation analysis of 57Fe after excess 57Fe sorption on clay showing precipitation of lepidocrocite. The quantitative repartition of mechanisms depicted in Figure 40 can be made (Figure 42) with the underlying hypothesis that no isotopic exchange between added 57Fe and clay structural Fe occurred in the course of their experiments. This hypothesis is, however, challenged by recent investigations (Handler et al. 2009; Beard et al. 2010; Rosso et al. 2010).

51

Figure 41. 56Fe sorption data from Schaefer et al. (2011).

Table 9. 56Fe sorption data from Schaefer et al. (2011)

Total Fe concentration

(mmol/L)

Equilibrium Fe concentration

(mmol/L)

Sorbed Fe concentration

(mmol/g)

0.91 0.0066 0.45 1.44 0.022 0.71 1.89 0.11 0.89 2.32 0.29 1.02 2.31 0.30 1.00 2.68 0.54 1.07

52

0.5 mmol/gclay57Fe (uptake > 99%) Natural 57Fe in clay structure

Clay Fe content × natural 57Fe abundance ratio4.54 × 2.119%

0.5 mmol/gclay57Fe(II)

Lepidocrocite Structural 57Fe(II) Sorbed 57Fe(III)

80%

17.5% = 0.105 mmol 57Fe/gclay

0.48 mmol Fe/gclay

No contribution from natural 57Fe in clay Structural Fe reduction

0.48 × 2.119% 0.01 mmol 57Fe/gclay

2.5% = 0.015 mmol 57Fe/gclay

+ 0.1 mmol/gclay57Fe(III)

Sorbed 57Fe(II)

+ 0.005 mmol 57Fe/gclay

Structural 57Fe(III)

Structural Fe reduction < 0.02 × 2.119%

0.09 mmol 57Fe/gclay0.015 mmol 57Fe/gclay

Figure 42. Quantitative analysis of 57Fe sorption data from Schaefer et al. (2011). In a first modelling step, we defined the simplest model to take into account the reaction scheme shown in Figure 40. The redox properties of structural iron were modelled assuming the following reaction: >Fe3+ + e- ↔ >Fe2+ log Kredox Reaction 6

Where >Fe stands for a structural Fe site. The activities of reduced and oxidized >Fe were assumed to be equal to their abundance relative to the total number of >Fe sites in a formalism similar to the ideal solid solution one:

3

2

redox K logFe

Felogpe Equation 7

Schaefer et al. (2011) could evidence that only 15% of structural Fe in nontronite was reducible by aqueous ferrous iron at pH 7.5. Hence, 68032 .FeFe mmol/g. Their experiments were conducted in an anaerobic glove-box but initial nontronite material was representative of oxidative conditions (100 % of >Fe3+). The initial redox of the clay suspension was not measured prior to the addition of solute ferrous iron but initial redox potential must have been positive to lead to lepidocrocite precipitation (Figure 43).

53

Figure 43. Predominance diagram of Fe in the experimental conditions of Schaefer et al. (2011) (in absence of clay). The diagram was calculated with PHREEPLOT (http://www.phreeplot.org/) using BRGM database (http://thermoddem.brgm.fr/). We modified our RSM model (Table 8) in order to reproduce the data from Schaefer et al. (2011):

the total strong site density was increased to 15 10-3 mol/kg (instead of 5 10-3) in order to be in agreement with the amount of sorbed Fe(III) calculated from Schaefer et al. data (Figure 42);

we added Reaction 6 and tried to adjust the value of log Kredox in order to reproduce data from Schaefer et al. Once log Kredox was chosen, initial pe of the

solution was set to a value that led to totFe%Fe 993 , i.e. initial pe > log

Kredox + 2 Once this condition was fulfilled, different values of initial pe led to very similar results;

The results are shown in Figure 44: the data are very satisfactorily reproduced with a log Kredox value between -1 and -2. This range of values corresponds to a structural iron standard redox potential of -120 to -60 mV, corresponding to a final Fe(II)/Fe(III) ratio of 0.05 to 0.15, based on total Fe concentration in clay. Note that only part of the structural iron participated in the reaction. A large part (85 %) of structural Fe did not react and should thus exhibit an even lower redox potential. This range of redox potential is in good agreement with very recent results reported by Gorski et al. (2012a,b) that were obtained through electrochemical techniques.

54

0.5 1 1.5 2 2.5 3

0.001

0.01

0.1

1

Added Fe(II) (mmol/L)

Eq

uili

briu

m 5

6F

e(II)

(m

mo

l/L)

log Kredox = -1

log Kredox = -2

log Kredox = -3

Exp. Data -1 -2 -30

0.1

0.4

0.5

log Kredox

"So

rbed

" 5

7F

e(II)

(m

mo

l/gcl

ay) Lepidocrocite

Sorbed Fe(III)sorbed Fe(II)

Figure 44. Comparison of RSM model + redox term with data from Schaefer et al. (2011). 3.6 Modelling of Fe(II) sorption on natural montmorillonite 3.6.1 MXBRGM The total structural Fe content in MX80 is ~0.45 mmol/gclay (ignoring the presence of oxides). In a first approach, we considered that all this iron had the same redox properties as those derived above for reactive iron in nontronite. We tried to predict sorption behaviour of MXBRGM first. We considered a log Kredox value of -1 in agreement with the modelling of data from Schaefer et al. (2011) (higher range value, see above). The initial Fe(II)/Fe(III) ratio in the structure was taken at the value measured by Mössbauer spectrometry, i.e. 0.1, and initial pe was adjusted accordingly. We considered an initial Fe sorbed amount of 7.4 10-6 mol/g as explained above and total mobile Fe was taken at the value indicated in Figure 30. The strong site capacity was varied to match the experimental data. This strong site capacity is necessarily above the amount of initial sorbed Fe, which does not correspond to the total site density (made of sites saturated with iron but also from ≡SOH or ≡SO- sites). Figure 45 shows that data at 0.3 and 0.5M NaCl are well predicted with a site density value of 8.2 mmol/kgclay, i.e. a total site density slightly above the total amount of mobile Fe (Table 10).

55

Figure 45. Comparison of predictive model for MXBRGM with available data at 0.3M NaCl and 0.05M NaCl. Fe solute concentrations are corrected values (see above)

Table 10. Parameters of MXBRGM predictive model (BPM; non electrostatic model).

Cation Exchange site density (X-) 0.79 mol kg-1 Cation Exchange reactions log K 2 XNa + Fe2+ ↔ X2Fe + 2 Na+ 0.5 XNa + FeCl+ ↔ XFeCl + Na+ 2 2 XNa + Ca2+ ↔ X2Ca + 2 Na+ 0.5 XNa + CaCl+ ↔ XCaCl + Na+ 2 Strong sites density 8.2 10-3 mol kg-1 Strong sites surface reactions log K ≡SOH ↔ ≡SO- + H+ -7.9 ≡SOH + Fe2+ ↔ ≡SOHFe2+ 4.5 ≡SOH + Fe3+ ↔ ≡SOFe2+ + H+ 16 ≡SOH + Fe3+ + H2O ↔ ≡SOFeOH+ + 2 H+ 11 ≡SOH + Fe3+ + 3 H2O ↔ ≡SOFe(OH)3

- + 4H+ -4 Weak sites density 8 10-2 mol kg-1 Weak sites surface reactions log K ≡WOH ↔ ≡WO- + H+ -7.9 ≡WOH + H+ ↔ ≡WOH2

+ 4.5 ≡WOH + Fe2+ + 2 H2O ↔ ≡WOFe(OH)2

- + 3 H+ -19 Redox sites density 0.45 mol kg-1 Redox sites reaction log K >Fe3+ + e- ↔ Fe2+ -1

56

3.6.2 MXVTT The initial Fe(II)/Fe(III) ratio in the structure was taken at the value measured by Mössbauer spectrometry, i.e. 0.3, and initial pe was adjusted accordingly. We considered an initial Fe sorbed amount of 1.9 10-6 mol/g as explained above and total mobile Fe was taken at the value indicated in Figure 30. Several hypotheses were tested to explain the lower affinity of Fe(II) for MXVTT compared to MXBRGM. The first hypothesis was to lower the strong site concentration at the clay surface, other parameters remaining identical to those tabulated in Table 10. This approach failed as it was not possible to reproduce Fe sorption curves on MXVTT by changing only this parameter. In a second approach we kept the parameters identical to those given in Table 10, excepting that the redox potential of structural Fe was changed. It was possible to fit satisfactorily the pH sorption edge data (A1 & A3) but not the Fe sorption isotherm as a function of total Fe concentration (note that sorption isotherm results is consistent with pH sorption edge for MXVTT whereas it was not the case for MXBRGM). Finally, we varied both parameters in order to find an acceptable fit of the data. Figure 46 (strong site capacity = 1.8 mmol/kgclay and log Kredox = -2.5) shows that a good agreement between model and data could be found but log Kd at pH>6 were still underestimated. We finally added the possibility of Fe(III)(OH)2

+ sorption on weak sites (log K = 2) in order to reproduce the increase in sorption at pH>6 (Figure 47). In this case, however, sorption extent prediction is far too high for pH > 7. If the same parameter is re-injected in the model for MXBRGM, sorption extent prediction is also too high compared to measured values (not shown).

2 4 6 8 100

2

4

6

8

pH

log

Kd

(K

d in

L/k

g)

Exp. data A1, centrifugationModel A1Exp. data A3, ultrafilteringModel A3

MXVTT, 0.3 M NaCl

2 4 6 8 102

3

4

5

6

7

8

9

pH

Exp. data B1, centrifugationModel B1Exp. data B3, ultrafilteringModel B3

MXVTT, 0.05 M NaCl

0 2x10-6 4x10-6 6x10-6 8x10-6 10-52.2

2.3

2.4

2.5

2.6

2.7

Fe equilibrium concentration (mol/L)

log

Kd

(K

d in

L/k

g)

Exp. data C1Model C1

MXVTT, 0.3 M NaCl, pH 5

Figure 46. Comparison of MXVTT sorption edge and isotherm data with MXVTT Fe sorption model (strong site capacity = 1.8 mmol/kgclay and log Kredox = -2.5).

2 4 6 8 102

4

6

8

10

12

pH

log

Kd

(K

d in

L/k

g)

Exp. data A1, centrifugationModel A1Exp. data A3, ultrafilteringModel A3

MXVTT, 0.3 M NaCl

2 4 6 8 102

4

6

8

10

12

pH

Exp. data B1, centrifugationModel B1Exp. data B3, ultrafilteringModel B3

MXVTT, 0.05 M NaCl

0 2x10-6 4x10-6 6x10-6 8x10-6 10-52.2

2.3

2.4

2.5

2.6

2.7

Fe equilibrium concentration (mol/L)

log

Kd

(K

d in

L/k

g)

Exp. data C1Model C1

MXVTT, 0.3 M NaCl, pH 5

Figure 47. Comparison of MXVTT sorption edge and isotherm data with MXVTT Fe sorption model (strong site capacity = 1.8 mmol/kgclay and log Kredox = -2.5, log K Fe(OH)2

+ on weak site = 2).

57

3.6.3 from Ph.D. thesis of Tournassat (2003) The BPM model was applied to data from the Ph.D. thesis of Tournassat (2003) without parameters adjustment (Table 10). The data are reasonably well predicted as long as the Fe surface loading does not exceed the strong site capacity (not shown). On the one hand, this blind prediction gives confidence on the model parameters for strong sites. On the other hand, these data evidence that other types of sites also participate in overall sorption for Fe surface occupancy greater than strong site capacity. By adjusting a log K value for Fe(III) sorption on weak sites, it is possible to reproduce the data correctly. However, (i) reintroducing this reaction in the BPM model leads to a strong overestimation of sorption results for MXBRGM and (ii) this value is not in agreement with the value obtained for MXVTT . 3.6.4 Summary of modelling results and challenges in finding a general model

for Fe(II) sorption on natural montmorillonites Table 11 summarizes the modelling constraints derived from above analysis. The underlying problem is that each reaction addition in the reaction scheme has an impact on the parameters that have been deduced from a previous modelling attempt. In the following, we propose a general model for the investigated clays by considering the following hypotheses:

Fe sorption is due to adsorption on exchange sites, strong and weak complexation sites and electron transfer with the structural Fe;

A non-electrostatic model derived from Bradbury and Baeyens (1997) can be used to model the data;

All mechanisms identified apply to all clay samples but with possible variations in (i) CEC values, (ii) structural Fe redox potential, and (ii) strong and weak sites surface density.

All these hypotheses could be challenged but we decided to stick to them in order to obtain a kind of “operational model”. 3.6.5 Proposition of a general model for montmorillonites and other smectites Modelling parameters according to the constraints from Table 11 as well as from data from Jaisi et al. (2008b) and modelling results are depicted in Figure 48. Jaisi et al. (2008b) performed Fe sorption experiments on nontronite as a function of pH, Fe concentrations and ionic strength. Their added Fe / clay ratio was high: 0.25 mol/kg. They studied the effect of contact time on the sorption extent between the Fe2+ solution and the clay and found that increasing contact time led to a marked increase in sorption, showing that after a long equilibrium time processes additional to sorption (e.g. nucleation and structural incorporation) play a role in Fe retention in clay materials. Their results are shown in Figure 49. The level of sorption and the different sorption region are relatively well reproduced for a one-day reaction (Figure 49). The plateau (100 %) at high pH is due to the precipitation of ferric oxide (for example lepidocrocite or ferrihydrite). After one day, the model predicts a precipitation edge at lower pH than actually observed. However, the ferric precipitate might be not at equilibrium, explaining the discrepancy and the tendency observed at 5 days and more: the second sorption edge is shifted towards low pH value with time.

58

Table 11. Summary of experiments together with related mechanisms and parameters.

Experiment Identified mechanisms and parameters Fe(II) sorption on synthetic smectite (Gehin et al., 2007 and this study)

Fe(II) affinity for strong sites Fe(III) affinity for strong sites Fe(II) affinity for weak sites

Fe(II) sorption on nontronite (Schaefer et al., 2011)

Electron transfer from the structure to the sorbed species Apparent redox potential for structural Fe

Fe(II) sorption on MXBRGM at low added Fe concentration (this study)

Relative importance of dissolved vs. desorption of Fe at low pH Confirmation of sorption parameters derived from Fe(II) sorption on nontronite and synthetic smectite. Strong site capacity variations as a function of clay nature

Fe(II) sorption on MXVTT at low added Fe concentration (this study)

Impact of clay preparation on (i) Fe(II)/Fe(III) structural ratio (ii) redox potential of structural Fe (iii) strong site capacity variations as a

function of clay preparation Possible electron transfer on weak sites and related Fe(III) affinity for weak sites

Fe(II) sorption on MX80 at low to high added Fe concentration (Tournassat, 2003)

Contribution of weak sites on sorption and necessary electron transfer

59

Tab

le 1

2. S

umm

ary

of F

e so

rpti

on p

aram

eter

s fo

r in

vest

igat

ed c

lays

(pa

ram

eter

s in

gra

y ce

lls

indi

cate

tha

t th

ey a

re t

he s

ame

for

all

inve

stig

ated

cla

ys)

.

A-

Syn

thet

ic

smec

tite

(Geh

in e

t al.

2007

)

B-

Syn

thet

ic

smec

tite

(Thi

s st

udy)

C

- M

XB

RG

M

D-

MX

VT

T

E-

MX

80

( Tou

rnas

sat

2003

)

F-

Non

tron

ite

NA

u-2

(Jai

si e

t al

. 200

8a)

G-

Non

tron

ite

NA

u-2

(Sch

aefe

r et a

l. 20

11)

Cat

ion

Exc

han

ge

C

EC

(m

ol/k

g)

0.5

0.87

0.79

0.98

0.69

0.7

0.7

log

K F

e2+

0.5

0.5

0.5

0.5

0.5

0.5

0.5

log

K F

eCl+

2

22

22

22

log

K C

a2+

0.5

0.5

0.5

0.5

0.5

0.5

0.5

log

K C

aCl+

2

22

22

22

Log

K H

+

11

11

11

1

R

edo

x

Str

uctu

ral F

e (m

ol/k

g)

00

0.45

0.45

0.45

0.68

0.68

F

e(II)

/Fe(

III)

ratio

0

00.

10.

30.

10.

010.

01

log

K r

edox

-9

9-9

9-2

-2.7

-2-1

.5-2

S

ite

a (s

tro

ng

sit

e)

S

ite d

ensi

ty (

mol

/kg)

0.

004

0.00

40.

0082

0.00

20.

008

0.02

0.02

lo

g K

sor

ptio

n F

e(II)

5

55

55

55

log

K s

orpt

ion

FeO

H+

-1-1

-1-1

-1-1

-1

log

K s

orpt

ion

Fe

3+

1616

1616

1616

16

log

K s

orpt

ion

FeO

H2+

10

.510

.510

.510

.510

.510

.510

.5

log

K s

orpt

ion

Fe(

OH

) 2+

3.5

3.5

3.5

3.5

3.5

3.5

3.5

Sit

e b

(w

eak

site

)

Site

den

sity

(m

ol/k

g)

0.08

0.08

0.08

0.08

0.1

0.15

0.15

lo

g so

rptio

n F

e(O

H) 2

-1

9.5

-19.

5-1

9.5

-19.

5-1

9.5

-19.

5-1

9.5

log

K s

orpt

ion

Fe

3+

1414

1414

1414

14

log

K s

orpt

ion

FeO

H2+

8

88

88

88

59

60

2 4 6 8 100

1

2

3

4

5

log

Kd

(K

d in

L/k

g)

2 4 6 8 1001234567

B B

2 4 6 8 100

1

2

3

4

5

pH

log

Kd

(K

d in

L/k

g)

2 4 6 8 100

1

2

3

4

5

6

pH

B B

2 4 6 8 100

20

40

60

80

100S

orb

ed F

e (%

)

2 4 6 8 100

10

20

30

40

50

60

% s

orb

ed F

e(III

)A A

Figure 48. Comparison of Fe sorption data (open symbols) with modelling predictions (closed symbols) for different smectites: A: Synthetic smectite (Gehin et al. 2007); B: synthetic smectite (this study); C: MXBRGM (this study, circles: centrifugation, triangles: ultrafiltering); D: MXVTT (this study, circles: centrifugation, triangles: ultrafiltering); E: MX80 (Tournassat, 2003); F: nontronite NAu-2 (Jaisi et al., 2008a); G: nontronite NAu-2 (Schaefer et al., 2011).

61

2 4 6 8 100

2

4

6

8lo

g K

d (

Kd

in L

/kg

)

2 4 6 8 100

2

4

6

8C D

2 4 6 8 100

2

4

6

8

10

log

Kd

(K

d in

L/k

g)

2 4 6 8 100

2

4

6

8

10C D

2 4 6 8 100

2

4

6

8

log

Kd

(K

d in

L/k

g)

2 4 6 8 100

2

4

6

8

10C D

2 4 6 8 100

2

4

6

8

pH

log

Kd

(K

d in

L/k

g)

2 4 6 8 100

2

4

6

8

10

pH

C D

Figure 48 (following). Comparison of Fe sorption data (open symbols) with modelling predictions (closed symbols) for different smectites: A: Synthetic smectite (Gehin et al., 2007); B: synthetic smectite (this study); C: MXBRGM (this study, circles: centrifugation, triangles: ultrafiltering); D: MXVTT (this study, circles: centrifugation, triangles: ultrafiltering); E: MX80 (Tournassat, 2003); F: nontronite NAu-2 (Jaisi et al. 2008a); G: nontronite NAu-2 (Schaefer et al. 2011).

62

2 4 6 8 100

0.002

0.004

0.006

0.008

0.01

0.012

pH

Sor

bed

Fe

(mol

/kg

)

2 4 6 8 10

0

0.005

0.01

0.015E

2 4 6 8 100

0.01

0.02

0.03

0.04

0.05

0.06E

2 4 6 8 100

0.020.040.060.08

0.10.120.14

E

2 4 6 8 100

0.010.020.030.040.050.060.07

pH

E E

2 4 6 8 100

0.001

0.002

0.003

0.004

0.005S

orb

ed F

e (m

ol/k

g) E

2 4 6 8 100

0.02

0.04

0.06

0.08

Sor

bed

Fe

(mol

/kg

) E

2 4 6 8 100

0.01

0.02

0.03

0.04

0.05

Sor

bed

Fe

(mol

/kg

) E

Figure 48 (following). Comparison of Fe sorption data (open symbols) with modelling predictions (closed symbols) for different smectites: A: Synthetic smectite (Gehin et al. 2007); B: synthetic smectite (this study); C: MXBRGM (this study, circles: centrifugation, triangles: ultrafiltering); D: MXVTT (this study, circles: centrifugation, triangles: ultrafiltering); E: MX80 (Tournassat 2003); F: nontronite NAu-2 (Jaisi et al. 2008a); G: nontronite NAu-2 (Schaefer et al., 2011).

63

2 4 6 8 100

0.01

0.02

0.03

0.04E

2 4 6 8 100

0.005

0.01

0.015

0.02S

orb

ed F

e (m

ol/k

g) E

2 4 6 8 100

20

40

60

80

100

pH

Sor

bed

Fe

(%) F

2 4 6 8 100

20

40

60

80

100

pH

F

2 4 6 8 100

20

40

60

80

100

pH

Sor

bed

Fe

(%) F

0 0.5 1 1.5 2 2.5 310-3

10-2

10-1

1

101

Total added Fe concentration (mmol/L)

Eq

. Fe

conc

entr

atio

n (m

mol

/L)

G

Lepidocrocitesaturation

Figure 48 (following). Comparison of Fe sorption data (open symbols) with modelling predictions (closed symbols) for different smectites: A: Synthetic smectite (Gehin et al., 2007); B: synthetic smectite (this study); C: MXBRGM (this study, circles: centrifugation, triangles: ultrafiltering); D: MXVTT (this study, circles: centrifugation, triangles: ultrafiltering); E: MX80 (Tournassat, 2003); F: nontronite NAu-2 (Jaisi et al. 2008a); G: nontronite NAu-2 (Schaefer et al. 2011).

64

Figure 49. Data from Jaisi et al. (2008b). The lines indicate the results from the model of Jaisi et al. (2008b) and not the result of the model developed in the present report. 3.7 Application of other data from the literature Few Fe sorption data on clay minerals exist in the literature and even fewer data are relevant to test the present model. Schultz and Grundl (2000; 2004) studied Fe sorption on a nontronite clay sample, but unfortunately no indication of ionic strength is given by the authors. It seems that a “zero” ionic strength was used (no added salt) thus favouring cation exchange over specific sorption. Merola et al. (2006) also investigated Fe sorption on nontronite clays but did not perform sorption edge nor isotherms. They demonstrate that measurable charge transfer from solute Fe(II) towards structural Fe(III) occurred through Fe(II) sorption and that this charge transfer occurred significantly for pH > 5 (at pH 5, no signal is detected whereas a significant signal is observed at pH 7). Their results are in agreement with the present data model which shows a marked increase of Fe sorption and oxidation for pH > 5.

65

3.8 Conclusion on the applicability of the model We could demonstrate that a sorption/redox model with an external calibration of clay structural Fe redox potential makes it possible to predict Fe(II) sorption onto montmorillonite. We validated this approach by applying the same model to Fe(II) sorption data on nontronite gathered in the literature. However, the redox properties of clay structural Fe are not constant. They seem to depend on the clay structure but also on the clay “history” reflected for example in their Fe(II)/Fe(III) ratio. For instance, Fe(II) sorption on two MX80 samples that were conditioned in two different laboratories led to contrasted sorption results that in turn led to a difference of the redox parameter of the model. We apply the developed model in the following section for the simulation of Fe(II) diffusional transport in a clay plug. The parameters for MX80VTT are used since the experiments were carried out with this material. Future use of this model should be considered with care due to the high dependence of its parameters to the exact nature of the clay sample. 3.9 Modelling of the diffusion experiments The diffusion experiments were modelled with PHREEQC and considered the following input parameters:

Cell size: 0.1 µm Montmorillonite partial dry density: 1.45 kg dm-3 Porosity: 0.47 Dp water: 1.2 10-10 m2/s

The model was run at two pH values (5.5 and 8) using the BPM model with a standard redox potential value of -240 mV for structural iron. The porewater Fe concentration was set at 1.5 µmol/L for the experiment at pH 5.5 in agreement with measurement of total iron at equilibrium in the sorption experiments. For the diffusion experiment at pH 8, the concentration was set at 0.1 µmol/L, which is a value slightly above the detection limit in the sorption experiment, and at 0.001 µmol/L. At pH 8, the calculated diffusion curve matches fairly well the observation for an Fe equilibrium concentration of 0.1 µmol/L (compare Figure 24 and Figure 50). The modelled Kd was 5 105 L/kg in these simulation conditions. This value is lower than the Kd value measured and modelled in the batch experiment. This is due to the consideration of an equilibrium Fe concentration of 0.1 µmol/L which must be higher than the Fe concentration present during the batch sorption experiment. However, the long experiment duration and the very high solid-to-liquid ratio in the diffusion experiment could have led to this concentration level due to a limited dissolution of iron hydroxide (remaining oxides are usually present even in treated clay samples, Tournassat et al. 2004) or clay structural iron (in the batch experiment similar dissolution would have led to very low concentration due to the dilution effect). If an equilibrium Fe concentration of 0.001 µmol/L is considered, calculation predicts that Fe radiotracer should not be detected in the second slice (Figure 51). These calculations

66

show how the system is sensitive to equilibrium conditions. Unfortunately, pore Fe concentration is not a parameter that can be obtained in the experimental conditions. At pH 5.5, the reactive transport model predicts that Fe should diffuse far away from the injection (Figure 52) in disagreement with experimental results. The reason for this discrepancy is not yet understood. A possible explanation could be a slow redox/sorption process which does not appear in the short batch sorption measurements. This interpretation is supported by the work of Jaisi (2008b, see above Section 3.6) and could be due to a slow inter-conversion of Fe(II) and Fe(III) in the clay lattice and subsequent renewal of edge sorption/oxidation sites.

0 0.0002 0.0004 0.0006 0.0008 0.0010

2

4

6

8

10

12

14

16

18

Distance (m)

Tra

cer

con

cen

trat

ion

(ar

bitr

ary)

Figure 50. Diffusion model results for the experiment at pH 8 with Fe equilibrium concentration of 0.1 µmol/L. The tracer concentration activity is given in an arbitrary unit vs. volume(similar to cpm/cm3).

67

0 0.0002 0.0004 0.0006 0.0008 0.0010

2

4

6

8

10

12

14

16

18

Distance (m)

Tra

cer

con

cen

trat

ion

(ar

bitr

ary)

Figure 51. Diffusion model results for the experiment at pH 8 with Fe equilibrium concentration of 0.001 µmol/L. The tracer concentration activity is given in an arbitrary unit vs. volume(similar to cpm/cm3).

0 0.0005 0.001 0.0015 0.0020

0.0001

0.0002

0.0003

0.0004

0.0005

0.0006

0.0007

Distance (m)

Tra

cer

con

cen

trat

ion

(ar

bitr

ary)

Figure 52. Diffusion model results for the experiment at pH 5.5. The tracer concentration activity is given in an arbitrary unit vs. volume(similar to cpm/cm3).

68

69

4 SUMMARY In this project, sorption and diffusion of Fe(II) in bentonite was studied. The experiments were carried out under low-oxygen conditions in an anaerobic glove-box. Radioactive isotope 55Fe was used as a tracer in the experiments. The sorption experiments were carried out with two batches of purified MX-80 bentonite. One was purified at BRGM (Bureau de Recherche Géologique et Minière, French Geological Survey) and the other one at VTT Finland. Experiments were also carried out with synthetic smectite, which did not include iron, prepared at LMPC (ENSC, F 68093 Mulhouse, France). The sorption experiments were carried out in 0.3 M and 0.05 M NaCl solutions as a function of pH, and in 0.3 M NaCl solution buffered at pH 5 as a function of added Fe(II) concentration. Separation of bentonite and solution at the end of the sorption experiment was carried out in the earlier phase by centrifuging only. Since there was an indication that some colloidal bentonite may stay in the solution, the separation method was improved in the later phase of the study. The supernatants from the centrifuging were then ultrafiltered in order to remove the colloidal bentonite from the solution. It appeared that the earlier studies had to be omitted when more than 97% of the tracer had been sorbed. Inactive iron is released from the bentonite to the solution in the sorption experiments. The natural iron originates from sorption sites and also from the dissolution of the clay network and/or the dissolution of accessory minerals. Qualitatively, the sorption of the radioactive 55Fe on all the clays shows the same type of behaviour, i.e. sorption increases with increasing pH. In the measurements with the bentonite purified at VTT, the sorption occurs at a higher pH than in the measurements carried out with bentonite purified at BRGM. The sorption experiments in the acetate buffer of pH 5 show decreasing sorption of 55Fe as a function of the increasing concentration of the added Fe(II). A general model for the investigated clays is proposed where Fe sorption is due to adsorption on exchange sites, strong and weak complexation sites and electron transfer with the structural Fe. All the mechanisms identified apply to all clay samples but with possible variations in CEC values, structural Fe redox potential and strong and weak sites' surface density. The diffusion experiments were carried out in compacted samples prepared from MXVTT and saturated with 0.3 M NaCl at pH 8 and 5. A non-steady-state diffusion experiment method, where the tracer is introduced as an impulse source between two bentonite plugs was used in the measurements. The measured diffusivities show rather low values (10-15 – 10-16 m2/s). At pH 8, the diffusion curve calculated with a reactive transport model on the basis of the sorption results match fairly well the experimental results. At pH 5, the model predicts much longer diffusion distance than found in the experiment, however. The reason for this discrepancy is not yet understood. A possible explanation could be a slow redox/sorption process which does not appear in the short batch sorption measurements.

70

71

REFERENCES Amman, L., Bergaya, F., Lagaly, G. 2005. Determination of the cation exchange capacity of clays with copper complexes revisited. Clay Minerals, vol. 40, 441 – 453. Baeyens, B., Bradbury, M. 1997. A mechanistic description of Ni and Zn sorption on Na-montmorillonite. Part I: Titration and sorption measurements. Journal of Contaminant Hydrology 27, 199-222. Beard, B. L., Handler, R. M., Scherer, M. M., Wu, L. L., Czaja, A. D., Heimann, A., and Johnson, C. M., 2010. Iron isotope fractionation between aqueous ferrous iron and goethite. Earth and Planetary Science Letters 295, 241-250. Bradbury, M., Baeyens, B. 1997. A mechanistic description of Ni and Zn sorption on Na-montmorillonite. Part II: Modelling. Journal of Contaminant Hydrology 27, 223-248. Bradbury, M., Baeyens, B. 2005. Modelling the sorption of Mn(II), Co(II), Ni(II), Zn(II), Cd(II), Eu(III), Am(III), Sn(IV), Th(IV), Np(V), and U(VI) on montmorillonite: Linear free energy relationships and estimates of surface binding constants for some selected heavy metals and actinides. Geochimica et Cosmochimica Acta 69, 875-892. Carlson, L., Karnland, O., Oversby, V., Rance, A., Smart, N., Snellman, M., Vähänen, M., Werme, L. 2007. Experimental studies of the interactions between anaerobically corroding iron and bentonite. Physics and Chemistry of the Earth 32, 334-345. Charlet, L., Sheinost, A.C., Tournassat, C., Greneche, J.M., Géhin, A., Fernández-Martinéz, A., Coudert, S., Tisserand, D., Brendle, J. 2007. Electron transfer at the mineral/water interface: Selenium reduction by ferrous iron sorbed on clay. Geochimica and Cosmochimica Acta, vol. 71, p. 5732-5749. Gehin, A., Greneche, J.-M., Tournassat, C., Brendle, J., Rancourt, D. G., and Charlet, L., 2007. Reversible surface-sorption-induced electron-transfer oxidation of Fe(II) at reactive sites on a synthetic clay mineral. Geochimica and Cosmochimica Acta, vol. 71, 863-876. Gorski, C. A., Aeschbacher, M., Soltermann, D., Voegelin, A., Baeyens, B., Marques Fernandes, M., Hofstetter, T. B., and Sander, M., 2012a. Redox Properties of Structural Fe in Clay Minerals. 1. Electrochemical Quantification of Electron-Donating and -Accepting Capacities of Smectites. Environ. Sci. Technol. 46, 9360-9368. Gorski, C. A., Klüpfel, L., Voegelin, A., Sander, M., and Hofstetter, T. B., 2012b. Redox Properties of Structural Fe in Clay Minerals. 2. Electrochemical and Spectroscopic Characterization of Electron Transfer Irreversibility in Ferruginous Smectite, SWa-1. Environ. Sci. Technol. 46, 9369-9377.

72

Handler, R. M., Beard, B. L., Johnson, C. M., and Scherer, M. M., 2009. Atom exchange between aqueous Fe(II) and goethite: An Fe isotope tracer study. Environ. Sci. Technol. vol. 43, p. 1102-1107. Hofstetter, T. B., Neumann, A., and Schwarzenbach, R. P., 2006. Reduction of nitroaromatic compounds by Fe(II) species associated with iron-rich smectites. Environ. Sci. Technol., vol. 40, p. 235-242. Hunter. F., Bate, F., Heath, T., Hoch, A. 2007. Geochemical investigation of iron transport into bentonite as steel corrodes. SKB TR-07-09. Idemitsu, K., Furuya, H., Inagagi, Y. 1993. Diffusion of corrosion products of iron in compacted bentonite. Mat. Res. Soc. Symp. Proc., vol. 294, p. 467-474. Idemitsu, K., Yano, S., Xiaopin, X., Inagagi, Y., Arima, T. 2002. Diffusion behavior of iron corrosion products in buffer materials. Mat. Res. Soc. Symp. Proc., vol. 713, p. 113-120. Jaisi, D. P., Dong, H. L., and Morton, J. P., 2008a. Partitioning of Fe(II) in reduced nontronite (NAu-2) to reactive sites: Reactivity in terms of Tc(VII) reduction. Clay. Clay. Miner., vol. 56, p. 175-189. Jaisi, D. P., Liu, C., Dong, H., Blake, R. E., and Fein, J. B., 2008b. Fe2+ sorption onto nontronite (NAu-2). Geochimica and Cosmochimica Acta, vol. 72, p. 5361-5371. Kamei, G., Oda, C., Mitsui, S., Shibata, S., Shinozaki, T. 1999. Fe(II)-Na ion exchange at interlayer of smectite: adsorption-desorption experiments and natural analogue. Engineering Geology, 54, 15-20. Lerf, A., Wagner, F. E., Poyato, J., and Perez-Rodriguez, J. L., 2011. Intercalation and dynamics of hydrated Fe2+ in the vermiculites from Santa Olalla and Ojen. J. Solid State Electrochem., vol. 15, p. 223-229. Merola, R. B., Fournier, E. D., and McGuire, M. M., 2006. Spectroscopic investigations of Fe2+ complexation on nontronite clay. Langmuir, vol. 23, p. 1223-1226. Muurinen, A., Carlsson, T. 2010. Experiences of pH and Eh measurements in compacted MX-80 bentonite. Applied Clay Science 47 (2010) 23 – 27. Sato, H., Ashida, T., Kohara, Y., Yui, M., Sasaki, N. 1992. Effect of dry density on diffusion of some radionuclides in compacted sodium bentonite. Journal of Nuclear Science and Technology, vol. 29, pp. 873 – 882. Rosso, K. M., Yanina, S. V., Gorski, C. A., Larese-Casanova, P., and Scherer, M. M., 2010. Connecting observations of hematite (alpha-Fe2O3) growth catalyzed by Fe(II). Environ. Sci. Technol., vol. 44, p. 61-67.

73

Schaefer, M. V., Gorski, C. A., and Scherer, M. M., 2011. Spectroscopic evidence for interfacial Fe(II)-Fe(III) electron transfer in a clay mineral. Environ. Sci. Technol., vol. 45, p. 540-545. Schultz, C. and Grundl, T., 2004. pH dependence of ferrous sorption onto two smectite clays. Chemosphere, vol. 57, p. 1301-1306. Schultz, C. A. and Grundl, T. J., 2000. pH dependence on reduction rate of 4-Cl-nitrobenzene by Fe(II)/montmorillonite systems. Environ. Sci. Technol., vol. 34, p. 3641-3648. Tournassat, C. 2003. Cations – clays interactions: the Fe(II) case. Application to the problematic of the French deep nuclear repository field concept. Thesis, Université de Joseph Fourier Grenoble-I, France. Tournassat, C., Greneche, J. M., Tisserand, D., and Charlet, L., 2004. The titration of clay minerals. Part I. Discontinuous backtitration technique combined to CEC measurements. J. Colloid Interf. Sci., vol. 273, p. 224-233. Tournassat, C., Gailhanou, H., Crouzet, C., Braibant, G., Gautier, A., and Gaucher, E. C., 2009. Cation exchange selectivity coefficient values on smectite and mixed-layer illite/smectite minerals. Soil Sci. Soc. Am. J. 73, 928-942. Tributh, H., Lagaly, G.A. 1986. Aufbereitung und Identifizierung von Boden- und Lagerstättentonen. I. Aufbereitung der Proben im Labor. GIT-Fachzeitschrift für das Laboratorium 30, 524 – 529. Welch, S., Beard, B., Johnson, C., Braterman, P. 2003. Geochimica et Cosmochimica Acta, vol. 67, no. 22, pp. 4231 – 4250. Wersin, P., Snellman, M., 2008. Impact of iron on the performance of clay barriers in waste disposal systems. Report on the status of research and development. Posiva Oy, Olkiluoto, Finland. Working Report 2008 – 7. Yao, S. Wang, M., and Madou, M., 2001. A pH electrode based on melt-oxidized iridium oxide. Journal of the Electrochemical Society, 148, (4) H29-H36. Yu, J-W., Neretnieks, I. 1997. Diffusion and sorption properties of radionuclides in compacted bentonite. SKB, Stockholm, Sweden. Technical Report 97-12.

74

A

PP

EN

DIX

A/1

S

orp

tion

exp

erim

ents

in 0

.3 M

NaC

l at

vary

ing

pH

wit

h M

X-b

ento

nit

e p

uri

fied

at

BR

GM

an

d V

TT

. S

epar

atio

n of

ben

toni

te f

rom

the

solu

tion

was

car

ried

out

wit

h a

cent

rifu

ge. T

he in

suff

icie

nt s

epar

atio

n co

ndit

ions

are

indi

cate

d w

ith

grey

ba

ckgr

ound

. The

init

ial F

e co

ncen

trat

ion

in th

e so

luti

on w

as 1

µm

ol/l

and

Fe-

55 a

ctiv

ity

613

823

cpm

.

Ben

ton

ite

sam

ple

F

e-55

tr

acer

so

luti

on

(m

l)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

MX

(BR

GM

)

MX

50

13.5

1.

5 10

0 H

Cl

2.13

69

5 98

8 51

7700

15

.7

1.9E

+02

M

X51

13

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1.5

10

HC

l 3.

21

397

441

3049

50

50.3

1.

0E+

03

MX

52

13.5

1.

5 0

--

4.20

10

8 14

4 14

2100

76

.9

3.3E

+03

M

X53

13

.5

1.5

1.0

NaO

H

4.46

61

.2

101

9555

0 84

.4

5.4E

+03

M

X54

13

.5

1.5

1.6

NaO

H

4.73

46

.5

73.9

46

300

92.5

1.

2E+

04

MX

55

13.5

1.

5 2.

0 N

aOH

5.

05

23.2

65

.5

3409

0 94

.4

1.7E

+04

M

X56

13

.5

1.5

2.2

NaO

H

5.63

16

.9

52.8

16

970

97.2

3.

5E+

04

MX

57

13.5

1.

5 2.

4 N

aOH

5.

81

19.0

65

.5

9750

98

.4

6.2E

+04

M

X58

13

.5

1.5

2.6

NaO

H

6.28

12

.7

29.6

96

65

98.4

6.

2E+

04

MX

59

13.5

1.

5 2.

8 N

aOH

6.

66

16.9

46

.5

5305

99

.1

1.1E

+05

M

X60

13

.5

1.5

3.0

NaO

H

6.96

19

.0

33.8

81

10

98.7

7.

5E+

04

M

X(V

TT

)

MX

60A

13

.5

1.5

100

HC

l 2.

13

532.

2 53

0.1

6063

50

1.2

1.2E

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M

X61

13

.5

1.5

10

HC

l 3.

19

361.

1 36

1.1

5813

50

5.3

5.6E

+01

M

X62

13

.5

1.5

0 --

4.

21

207.

0 20

4.8

5362

50

12.6

1.

4E+

02

MX

63

13.5

1.

5 0.

8 N

aOH

4.

52

160.

5 16

0.5

5121

00

16.6

2.

0E+

02

MX

64

13.5

1.

5 1.

2 N

aOH

4.

73

126.

7 13

3.0

4842

50

21.1

2.

7E+

02

MX

65

13.5

1.

5 1.

4 N

aOH

5.

04

111.

9 10

9.8

4328

50

29.5

4.

2E+

02

MX

66

13.5

1.

5 1.

6 N

aOH

5.

39

73.9

78

.1

3349

50

45.4

8.

3E+

02

MX

67

13.5

1.

5 1.

8 N

aOH

6.

12

19.0

25

.3

7216

0 88

.2

7.5E

+03

M

X68

13

.5

1.5

2.0

NaO

H

6.28

<

5 10

.6

7575

98

.8

8.0E

+04

M

X69

13

.5

1.5

2.2

NaO

H

6.72

10

.6

12.7

84

75

98.6

7.

1E+

04

MX

70

13.5

1.

5 2.

5 N

aOH

8.

11

8.4

14.8

13

735

97.8

4.

4E+

04

75

A

PP

EN

DIX

A/2

S

orp

tion

exp

erim

ents

in 0

.3 M

NaC

l at

vary

ing

pH

wit

h s

ynth

etic

sm

ecti

te (

SB

) p

rep

ared

at

BR

GM

an

d w

ith

out

ben

ton

ite

(CT

).

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

. The

insu

ffic

ient

sep

arat

ion

cond

itio

ns a

re in

dica

ted

wit

h gr

ey

back

grou

nd. T

he in

itia

l Fe

conc

entr

atio

n in

the

solu

tion

was

1 µ

mol

/l a

nd F

e-55

act

ivit

y 61

3 82

3 cp

m.

B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

Syn

thet

ic s

mec

tite

(S

B)

SB

51

13.5

1.

5 10

0 H

Cl

2.12

71

.8

71.8

60

7850

1.

0 9.

8E+

00

SB

52

13.5

1.

5 10

H

Cl

3.18

63

.4

57.0

56

6700

7.

7 8.

3E+

01

SB

53

13.5

1.

5 0.

0 --

4.

12

61.2

54

.9

5088

00

17.1

2.

1E+

02

SB

54

13.5

1.

5 1.

5 N

aOH

4.

92

59.1

50

.7

4608

50

24.9

3.

3E+

02

SB

55

13.5

1.

5 2.

0 N

aOH

5.

53

69.7

57

.0

3863

20

37.1

5.

9E+

02

SB

56

13.5

1.

5 2.

4 N

aOH

6.

50

33.8

29

.6

2664

00

56.6

1.

3E+

03

SB

57

13.5

1.

5 2.

5 N

aOH

6.

57

33.8

27

.5

2058

50

66.5

2.

0E+

03

SB

58

13.5

1.

5 2.

6 N

aOH

6.

72

21.1

10

.6

1702

85

72.3

2.

6E+

03

SB

59

13.5

1.

5 2.

8 N

aOH

7.

01

16.9

16

.9

9248

0 84

.9

5.6E

+03

S

B60

13

.5

1.5

3.2

NaO

H

7.60

8.

4 <

5 22

840

96.3

2.

6E+

04

No

ben

ton

ite

in t

he

cen

trif

ug

al t

ub

e (C

T)

C

T51

13

.5

1.5

100

HC

l 2.

24

52.8

57

.0

6201

50

-1.0

CT

52

13.5

1.

5 10

H

Cl

3.16

50

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50.7

61

7550

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C

T53

13

.5

1.5

0.0

--

3.86

54

.9

48.6

61

7100

-0

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C

T54

13

.5

1.5

1.0

NaO

H

4.80

52

.8

48.6

61

3000

0.

1

CT

55

13.5

1.

5 2.

2 N

aOH

5.

05

52.8

50

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6110

50

0.5

C

T56

13

.5

1.5

2.3

NaO

H

5.38

50

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46.5

60

6900

1.

1

CT

57

13.5

1.

5 2.

4 N

aOH

5.

94

54.9

46

.5

6124

00

0.2

C

T58

13

.5

1.5

2.5

NaO

H

7.38

50

.7

40.1

53

3750

13

.0

C

T59

13

.5

1.5

2.8

NaO

H

8.93

19

.0

25.3

30

9170

49

.6

C

T60

13

.5

1.5

3.2

NaO

H

9.42

12

.7

14.8

22

7955

62

.9

76

AP

PE

ND

IX A

/3

Sor

pti

on e

xper

imen

ts in

0.3

M N

aCl a

t va

ryin

g p

H

Exp

erim

ents

wit

h M

X-b

ento

nite

pur

ifie

d at

BR

GM

and

VT

T.

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

fol

low

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y ul

traf

iltr

atio

n.

Init

ial F

e co

ncen

trat

ion

in th

e so

luti

on w

as 1

.1 µ

mol

/l a

nd F

e-55

act

ivit

y 60

2 49

3 cp

m.

B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

MX

(BR

GM

)

M

XB

11

13.5

1.

5 0.

0 --

--

4.30

43

.9

41.8

85

806

85.8

6.

0E+

03

MX

B12

13

.5

1.5

0.5

NaO

H

4.49

27

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27.2

56

576

90.6

9.

6E+

03

MX

B13

13

.5

1.5

1.2

NaO

H

5.59

8.

4 14

.6

1198

1 98

.0

4.9E

+04

M

XB

14

13.5

1.

5 2.

2 N

aOH

6.

91

<5

<5

125

100.

0 4.

8E+

06

MX

B15

13

.5

1.5

3.5

NaO

H

8.89

<

5 <

5 co

ntam

in.

MX

B16

13

.5

1.5

4.5

NaO

H

9.39

<

5 <

5 19

9 10

0.0

3.0E

+06

M

XB

17

13.5

1.

5 2.

5 N

aOH

7.

63

<5

cont

amin

. 19

10

0.0

3.2E

+07

M

X(V

TT

)

M

XV

11

13.5

1.

5 0.

0 --

--

4.25

15

2.7

150.

6 45

1945

25

.0

3.3E

+02

M

XV

12

13.5

1.

5 0.

5 N

aOH

4.

75

108.

8 11

0.9

3946

97

34.5

5.

3E+

02

MX

V13

13

.5

1.5

0.8

NaO

H

5.40

64

.8

62.7

25

6379

57

.4

1.3E

+03

M

XV

14

13.5

1.

5 1.

2 N

aOH

6.

27

<5

<5

2114

99

.6

2.8E

+05

M

XV

15

13.5

1.

5 1.

7 N

aOH

7.

51

<5

<5

22

100.

0 2.

7E+

07

MX

V16

13

.5

1.5

2.2

NaO

H

8.59

<

5 <

558

10

0.0

1.0E

+07

M

XV

17

13.5

1.

5 1.

7 N

aOH

7.

94

<5

<5

26

100.

0 2.

4E+

07

77

A

PP

EN

DIX

A/4

S

orp

tion

exp

erim

ents

in 0

.3 M

NaC

l at

vary

ing

pH

E

xper

imen

ts w

ith

synt

heti

c sm

ecti

te (

SB

) pr

epar

ed a

t BR

GM

and

wit

hout

ben

toni

te (

CT

).

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

fol

low

ed b

y ul

traf

iltr

atio

n. T

he in

itia

l Fe

conc

entr

atio

n in

the

solu

tion

was

1.1

µm

ol/l

and

Fe-

55 a

ctiv

ity

602

493

cpm

.

Ben

ton

ite

sam

ple

F

e-55

tr

acer

so

luti

on

(m

l)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

Syn

thet

ic s

mec

tite

(S

B)

S

B11

13

.5

1.5

0.0

----

4.

35

48.1

48

.1

4433

74

26.4

3.

6E+

02

SB

12

13.5

1.

5 0.

5 N

aOH

4.

62

46.0

48

.1

4242

19

29.6

4.

2E+

02

SB

13

13.5

1.

5 1.

0 N

aOH

5.

71

43.9

46

.0

3818

92

36.6

5.

8E+

02

SB

14

13.5

1.

5 1.

8 N

aOH

6.

64

16.7

27

.2

1570

81

73.9

2.

8E+

03

SB

15

13.5

1.

5 2.

7 N

aOH

7.

98

<5

8.4

2738

99

.5

2.2E

+05

S

B16

13

.5

1.5

3.5

NaO

H

8.51

<

5 12

.5

253

100.

0 2.

4E+

06

SB

17

13.5

1.

5 2.

8 N

aOH

8.

05

<5

16.7

32

29

99.5

1.

9E+

05

No

ben

ton

ite

in t

he

cen

trif

ug

al t

ub

e (C

T)

C

T11

13

.5

1.5

0.0

----

3.

99

62.7

62

.7

6011

63

0.2

C

T12

13

.5

1.5

1.4

NaO

H

4.75

71

.1

60.7

60

3823

-0

.2

C

T13

13

.5

1.5

1.6

NaO

H

5.93

64

.8

69.0

59

3325

1.

5

CT

14

13.5

1.

5 1.

8 N

aOH

6.

70

54.4

56

.5

5187

41

13.9

CT

15

13.5

1.

5 1.

9 N

aOH

8.

30

31.4

37

.6

3183

22

47.2

CT

16

13.5

1.

5 2.

1 N

aOH

8.

85

16.7

20

.9

1733

88

71.2

78

A

PP

EN

DIX

B/1

S

orp

tion

exp

erim

ents

in 0

.05

M N

aCl a

t va

ryin

g p

H

Exp

erim

ents

wit

h M

X-b

ento

nite

pur

ifie

d at

BR

GM

and

VT

T.

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

. The

insu

ffic

ient

sep

arat

ion

cond

itio

ns a

re in

dica

ted

wit

h gr

ey

back

grou

nd. T

he in

itia

l Fe

conc

entr

atio

n in

the

solu

tion

was

1 µ

mol

/l a

nd F

e-55

act

ivit

y 61

3 82

3 cp

m.

B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

MX

(BR

GM

)

MX

71

13.5

1.

5 10

0 H

Cl

2.15

68

2 78

4 39

5100

35

.6

5.5E

+02

M

X72

13

.5

1.5

10

HC

l 3.

24

381

406

2682

00

56.3

1.

3E+

03

MX

73

13.5

1.

5 0.

0 --

4.

30

92

119

9675

0 84

.2

5.4E

+03

M

X74

13

.5

1.5

1.5

NaO

H

4.65

46

77

46

100

92.5

1.

2E+

04

MX

75

13.5

1.

5 2.

0 N

aOH

5.

05

46

56

1755

5 97

.1

3.4E

+04

M

X76

13

.5

1.5

2.4

NaO

H

5.78

21

65

99

70

98.4

6.

1E+

04

MX

77

13.5

1.

5 2.

5 N

aOH

6.

06

23

48

7495

98

.8

8.1E

+04

M

X78

13

.5

1.5

2.6

NaO

H

6.67

15

42

97

20

98.4

6.

2E+

04

MX

79

13.5

1.

5 2.

8 N

aOH

7.

03

23

63

1329

0 97

.8

4.5E

+04

M

X80

13

.5

1.5

3.2

NaO

H

7.31

33

56

13

385

97.8

4.

5E+

04

M

X(V

TT

)

MX

81

13.5

1.

5 10

0 H

Cl

2.16

49

4 49

2 51

5300

16

.1

1.9E

+02

M

X82

13

.5

1.5

10

HC

l 3.

21

389

316

4573

50

25.5

3.

4E+

02

MX

83

13.5

1.

5 0.

0 --

4.

24

165

161

3761

50

38.7

6.

3E+

02

MX

84

13.5

1.

5 1.

0 N

aOH

4.

66

111

111

3361

50

45.2

8.

3E+

02

MX

85

13.5

1.

5 2.

2 N

aOH

5.

01

88

86

2893

00

52.9

1.

1E+

03

MX

86

13.5

1.

5 2.

3 N

aOH

5.

32

69

69

2472

00

59.7

1.

5E+

03

MX

87

13.5

1.

5 2.

4 N

aOH

5.

89

27

27

7944

0 87

.1

6.7E

+03

M

X88

13

.5

1.5

2.5

NaO

H

6.28

15

8

8840

98

.6

6.8E

+04

M

X89

13

.5

1.5

2.8

NaO

H

6.88

15

38

19

380

96.8

3.

1E+

04

MX

90

13.5

1.

5 3.

2 N

aOH

8.

70

38

84

5115

0 91

.7

1.1E

+04

79

A

PP

EN

DIX

B/2

S

orp

tion

exp

erim

ents

in 0

.05

M N

aCl a

t va

ryin

g p

H

Exp

erim

ents

wit

h sy

nthe

tic

smec

tite

(S

B)

prep

ared

at B

RG

M a

nd w

itho

ut b

ento

nite

(C

T).

S

epar

atio

n of

ben

toni

te f

rom

the

solu

tion

was

car

ried

out

wit

h a

cent

rifu

ge. T

he in

suff

icie

nt s

epar

atio

n ha

s be

en in

dica

ted

wit

h th

e gr

ey

back

grou

nd. T

he in

itia

l Fe

conc

entr

atio

n in

the

solu

tion

was

1 µ

mol

/l a

nd F

e-55

act

ivit

y 61

3 82

3 cp

m.

B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

Syn

thet

ic s

mec

tite

(S

B)

SB

61

13.5

1.

5 10

0 H

Cl

2.12

co

ntam

in.

59

5044

50

17.8

2.

2E+

02

SB

62

13.5

1.

5 10

H

Cl

3.26

67

50

42

8100

30

.3

4.3E

+02

S

B63

13

.5

1.5

0 --

4.

30

44

52

3566

50

41.9

7.

2E+

02

SB

64

13.5

1.

5 1.

2 N

aOH

4.

74

50

33

3207

00

47.8

9.

1E+

02

SB

65

13.5

1.

5 1.

8 N

aOH

5.

18

42

29

2920

50

52.4

1.

1E+

03

SB

66

13.5

1.

5 2.

0 N

aOH

5.

70

44

33

2638

00

57.0

1.

3E+

03

SB

67

13.5

1.

5 2.

2 N

aOH

5.

95

23

25

2199

35

64.2

1.

8E+

03

SB

68

13.5

1.

5 2.

5 N

aOH

6.

61

21

17

1464

15

76.1

3.

2E+

03

SB

69

13.5

1.

5 2.

8 N

aOH

7.

02

25

15

7977

0 87

.0

6.7E

+03

S

B70

13

.5

1.5

3.2

NaO

H

7.80

21

13

24

775

96.0

2.

4E+

04

No

ben

ton

ite

in t

he

cen

trif

ug

al t

ub

e (C

T)

C

T61

13

.5

1.5

100

HC

l 2.

17

52

52

6096

50

0.7

C

T62

13

.5

1.5

10

HC

l 3.

08

52

50

6043

00

1.6

C

T63

13

.5

1.5

0 --

3.

79

50

50

6122

50

0.3

C

T64

13

.5

1.5

1.5

NaO

H

4.17

52

50

60

3450

1.

7

CT

65

13.5

1.

5 2.

0 N

aOH

4.

47

52

48

6050

00

1.4

C

T66

13

.5

1.5

2.2

NaO

H

4.69

52

48

60

8650

0.

8

CT

67

13.5

1.

5 2.

3 N

aOH

4.

90

52

48

6012

00

2.1

C

T68

13

.5

1.5

2.4

NaO

H

5.15

52

48

59

8250

2.

5

CT

69

13.5

1.

5 2.

5 N

aOH

5.

72

52

50

6021

00

1.9

C

T70

13

.5

1.5

2.6

NaO

H

8.45

36

40

45

7745

25

.4

80

AP

PE

ND

IX B

/3

Sor

pti

on e

xper

imen

ts in

0.0

5 M

NaC

l at

vary

ing

pH

E

xper

imen

ts w

ith

MX

-ben

toni

te p

urif

ied

at B

RG

M a

nd V

TT

. S

epar

atio

n of

ben

toni

te f

rom

the

solu

tion

was

car

ried

out

wit

h a

cent

rifu

ge f

ollo

wed

by

ultr

afil

trat

ion.

Ini

tial

Fe

conc

entr

atio

n in

the

solu

tion

was

1.1

µm

ol/l

and

Fe-

55 a

ctiv

ity

612

159

cpm

. B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

pH

at

the

end

of

the

test

Fe(

II) a

t th

e en

d o

f th

e te

st

(µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

MX

(BR

GM

)

M

XB

1 13

.5

1.5

0 --

--

4.59

27

.2

25.1

44

836

92.7

1.

3E+

04

M

XB

3 13

.5

1.5

1.5

NaO

H

6.41

<

5 co

ntam

in.

1056

99

.8

5.8E

+05

M

XB

4 13

.5

1.5

1.8

NaO

H

6.52

<

5 <

5 24

25

99.6

2.

5E+

05

MX

B5

13.5

1.

5 2.

1 N

aOH

6.

85

<5

<5

94

100.

0 6.

5E+

06

MX

B6

13.5

1.

5 2.

4 N

aOH

7.

07

<5

<5

81

100.

0 7.

6E+

06

MX

B7

13.5

1.

5 2.

1 N

aOH

7.

12

<5

<5

133

100.

0 4.

6E+

06

MX

(VT

T)

MX

V1

13.5

1.

5 0

----

4.

6 12

5.5

121.

3 32

3695

47

.1

8.1E

+02

M

XV

2 13

.5

1.5

0.6

NaO

H

5.5

73.2

71

.1

2479

42

59.5

1.

5E+

03

MX

V3

13.5

1.

5 1.

0 N

aOH

6.

5 <

5 <

512

346

98.0

4.

9E+

04

MX

V4

13.5

1.

5 1.

3 N

aOH

7.

04

<5

<5

65

100.

0 9.

5E+

06

MX

V5

13.5

1.

5 1.

6 N

aOH

7.

62

<5

<5

43

100.

0 1.

4E+

07

MX

V6

13.5

1.

5 1.

9 N

aOH

8.

34

<5

<5

cont

amin

.

M

XV

7 13

.5

1.5

1.9

NaO

H

8.6

<5

<5

41

100.

0 1.

5E+

07

81

AP

PE

ND

IX B

/4

Sor

pti

on e

xper

imen

ts in

0.0

5 M

NaC

l at

vary

ing

pH

E

xper

imen

ts w

ith

synt

heti

c sm

ecti

te (

SB

) pr

epar

ed a

t BR

GM

and

wit

hout

ben

toni

te (

CT

).

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

fol

low

ed b

y ul

traf

iltr

atio

n.

Init

ial F

e co

ncen

trat

ion

in th

e so

luti

on w

as 1

.1 µ

mol

/l a

nd F

e-55

act

ivit

y 61

2 15

9 cp

m.

B

ento

nit

e sa

mp

le

Fe-

55

trac

er

solu

tio

n

(ml)

Ben

ton

ite

dis

per

sio

n

(ml)

NaO

H o

r H

Cl a

dd

ed

for

pH

ad

just

men

t

(µm

ol)

Ad

ded

co

mp

ou

nd

p

H a

t th

e en

d o

f th

e te

st

Fe(

II) a

t th

e en

d o

f th

e te

st (

µg

/l)

Fe(

tot)

at

the

end

of

the

test

(µ

g/l)

Fe-

55

acti

vity

at

the

end

of

the

test

(c

pm

)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

Syn

thet

ic s

mec

tite

(S

B)

SB

1 13

.5

1.5

0 --

--

4.85

31

.4

29.3

28

7320

53

.1

1.1E

+03

S

B2

13.5

1.

5 1.

0 N

aOH

5.

92

20.9

18

.8

1905

24

68.9

2.

2E+

03

SB

3 13

.5

1.5

1.5

NaO

H

6.6

12.5

12

.5

1002

34

83.6

5.

1E+

03

SB

4 13

.5

1.5

1.8

NaO

H

6.9

6.3

12.5

57

085

90.7

9.

7E+

03

SB

5 13

.5

1.5

2.1

NaO

H

7.52

<

56.

3 14

687

97.6

4.

1E+

04

SB

6 13

.5

1.5

2.4

NaO

H

7.55

<

5<

519

170

96.9

3.

1E+

04

SB

7 13

.5

1.5

3.0

NaO

H

8.29

<

5<

586

2 99

.9

7.1E

+05

N

o b

ento

nit

e in

th

e ce

ntr

ifu

gal

tu

be

(CT

)

CT

1 13

.5

1.5

0 --

--

3.85

64

.8

62.7

60

8689

0.

6

CT

2 13

.5

1.5

1.0

NaO

H

4.29

64

.8

71.1

61

2801

-0

.1

C

T3

13.5

1.

5 1.

3 N

aOH

4.

55

66.9

62

.7

6136

84

-0.2

CT

4 13

.5

1.5

1.6

NaO

H

5.78

62

.7

60.7

61

3461

-0

.2

C

T5

13.5

1.

5 1.

8 N

aOH

7.

41

60.7

33

.5

3345

53

45.3

CT

6 13

.5

1.5

2.0

NaO

H

8.65

31

.4

29.3

27

0803

55

.8

82

AP

PE

ND

IX C

/1

Sor

pti

on e

xper

imen

ts in

ace

tate

bu

ffer

of

pH 5

at

vary

ing

init

ial i

ron

con

cen

trat

ion

s

Exp

erim

ents

wit

h M

X-b

ento

nite

pur

ifie

d at

BR

GM

and

VT

T.

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

fol

low

ed b

y ul

traf

iltr

atio

n.

Init

ial t

otal

act

ivit

y of

Fe-

55 w

as 6

88 7

25 c

pm.

S

amp

le

Init

ial i

ron

co

nce

ntr

atio

n

(µg

/l)

pH

at

the

end

F

inal

Fe(

II)

(µg

/l)

Fin

al F

e(to

t)

(µg

/l)

Fin

al F

e-55

ac

tivi

ty i

n

solu

tio

n

(cp

m)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

MX

(BR

GM

)

MX

91

45

4.94

0.

0 10

.5

2251

0 96

.7

3.0E

+04

M

X92

10

0 4.

95

2.1

18.8

25

745

96.3

2.

6E+

04

MX

93

156

4.95

4.

2 14

.6

3235

0 95

.3

2.0E

+04

M

X94

21

2 4.

95

16.7

29

.3

5015

5 92

.7

1.3E

+04

M

X95

26

8 4.

94

20.9

33

.5

5499

0 92

.0

1.2E

+04

M

X96

32

4 4.

95

41.8

54

.4

9744

0 85

.9

6.1E

+03

M

X97

37

9 4.

93

50.2

10

2.5

4520

0 93

.4

1.4E

+04

M

X98

43

5 4.

93

79.5

10

2.5

9838

0 85

.7

6.0E

+03

M

X99

49

1 4.

93

135.

9 96

.2

8614

5 87

.5

7.0E

+03

M

X10

0 54

7 4.

93

98.3

11

2.9

9641

5 86

.0

6.1E

+03

M

X(V

TT

)

MX

101

45

4.94

co

ntam

in.

cont

amin

. 43

3900

37

.0

5.9E

+02

M

X10

2 10

0 4.

95

131.

8 14

6.4

4727

50

31.4

4.

6E+

02

MX

103

156

4.93

17

9.9

207.

1 50

5550

26

.6

3.6E

+02

M

X10

4 21

2 4.

93

246.

8 26

5.6

5242

00

23.9

3.

1E+

02

MX

105

268

4.93

28

2.4

305.

4 54

1000

21

.4

2.7E

+02

M

X10

6 32

4 4.

93

332.

6 34

9.3

5427

00

21.2

2.

7E+

02

MX

107

379

4.93

38

2.7

395.

3 56

0250

18

.7

2.3E

+02

M

X10

8 43

5 4.

94

432.

9 46

8.5

5667

00

17.7

2.

1E+

02

MX

109

491

4.99

51

0.3

510.

3 57

5600

16

.4

2.0E

+02

M

X11

0 54

7 5.

00

560.

5 55

8.4

5842

50

15.2

1.

8E+

02

83

A

PP

EN

DIX

C/2

S

orp

tion

exp

erim

ents

in a

ceta

te b

uff

er o

f pH

5 a

t va

ryin

g in

itia

l iro

n c

once

ntr

atio

ns

E

xper

imen

ts w

ith

synt

heti

c sm

ecti

te (

SB

) pr

epar

ed a

t BR

GM

and

wit

hout

ben

toni

te (

CT

).

Sep

arat

ion

of b

ento

nite

fro

m th

e so

luti

on w

as c

arri

ed o

ut w

ith

a ce

ntri

fuge

fol

low

ed b

y ul

traf

iltr

atio

n.

Init

ial t

otal

act

ivit

y of

Fe-

55 w

as 6

88 7

25 c

pm.

S

amp

le

Init

ial i

ron

co

nce

ntr

atio

n

(µg

/l)

pH

at

the

end

F

inal

Fe(

II)

(µg

/l)

Fin

al F

e(to

t)

(µg

/l)

Fin

al F

e-55

ac

tivi

ty i

n

solu

tio

n

(cp

m)

So

rbed

of

Fe-

55 (

%)

Kd (

ml/g

)

SB

(BR

GM

)

SB

71

45

5.00

63

65

47

7810

30

.6

4.4E

+02

S

B72

10

0 5.

00

cont

amin

atio

n88

49

8914

27

.6

3.8E

+02

S

B73

15

6 5.

00

140

151

5105

52

25.9

3.

5E+

02

SB

74

212

5.00

18

2 18

4 52

5379

23

.7

3.1E

+02

S

B75

26

8 4.

98

cont

amin

atio

n26

8 54

0328

21

.5

2.7E

+02

S

B76

32

4 4.

98

276

305

5818

50

15.5

1.

8E+

02

SB

77

379

4.99

32

6 33

3 57

9550

15

.9

1.9E

+02

S

B78

43

5 4.

98

389

383

5926

50

13.9

1.

6E+

02

SB

79

491

4.99

43

1 43

7 59

3650

13

.8

1.6E

+02

S

B80

54

7 4.

99

485

489

5978

50

13.2

1.

5E+

02

No

ben

ton

ite

in t

he

cen

trif

ug

al t

ub

e (C

T)

C

T71

45

4.

97

44

73

6777

50

1.6

C

T72

10

0 4.

98

105

117

6906

50

-0.3

CT

73

156

4.99

24

9 16

7 68

0500

1.

2

CT

74

212

4.99

21

3 22

4 69

6600

-1

.1

C

T75

26

8 4.

99

276

282

6915

50

-0.4

CT

76

324

4.98

32

8 33

3 69

0400

-0

.2

C

T77

37

9 4.

98

310

619

6890

50

0.0

C

T78

43

5 4.

98

356

443

6923

50

-0.5

CT

79

491

4.98

49

4 50

2 68

7250

0.

2

CT

80

547

4.98

55

0 55

4 69

1150

-0

.4

84