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Dissolution-recrystallization of (Mg,Fe)CO during hydrothermal cycles: Fe/Fe conundrums in the carbonation of ferromagnesian minerals. II

III

Fulvio Di Lorenzo, and Manuel Prieto Cryst. Growth Des., Just Accepted Manuscript • DOI: 10.1021/acs.cgd.7b00438 • Publication Date (Web): 23 Jun 2017 Downloaded from http://pubs.acs.org on June 27, 2017

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Crystal Growth & Design

Dissolution-recrystallization of (Mg,Fe)CO3 during hydrothermal cycles: FeII/FeIII conundrums in the carbonation of ferromagnesian minerals Fulvio Di Lorenzo and Manuel Prieto Department of Geology, C/Jesús Arias de Velasco s/n, University of Oviedo, 33005 Oviedo, Spain Cover page This study highlights some factors that can control the artificial carbonation of basaltic rocks for CO2 sequestration purposes, such as the influence of redox processes on the crystallization behavior of ferromagnesian carbonates. Crystallization experiments were performed at high supersaturation with the aim of obtaining metastable (Mg,Fe)CO3 precipitates to study their evolution during hydrothermal heating (150-180 ºC and 5-9 bars) in a nitrogen atmosphere. The observation of crystals with different compositions show that, during the heating period, the primary (Mg,Fe)CO3 precipitates undergo a dissolution-recrystallization process driven by two forces: The preferential partitioning of the aqueous Fe2+ ions towards the (Mg,Fe)CO3 solid solution and the oxidation of aqueous Fe2+ species to precipitate Fe(III)-bearing solids. The gradual oxidation of the aqueous Fe2+ species eventually leads to the crystallization of magnetite and amorphous Fe(III)-bearing precipitates. Magnetite is also metastable since the aqueous solution is maximally supersaturated with respect to hematite, at least if we assume that the equilibrium distribution of aqueous iron species is attained instantaneously. However, reactions between aqueous species are actually not instantaneous and can be slower than the dissolution/precipitation of related solids. Considering the coupling between oxidation and precipitation of FeCO3 is essential in modeling the artificial carbonation of ferromagnesian minerals. Abstract Crystallization experiments were performed at high supersaturation to obtain metastable (Mg,Fe)CO3 precipitates and study their evolution during hydrothermal heating in a nitrogen atmosphere. The primary (Mg,Fe)CO3 solids undergo a dissolution-recrystallization process driven by two forces: The preferential partition of Fe2+ towards the (Mg,Fe)CO3 solid solution and the oxidation of aqueous Fe2+ species to form Fe(III)-bearing solids. Under limited oxygen supply, the oxidation of Fe2+ is a slow process, particularly in the presence of chlorine and inorganic carbon. In absence of oxidation, the initial solid solution should become increasingly Fe-rich as the supersaturation decreases and the system approaches equilibrium. Here, the secondary (Mg,Fe)CO3 solids become the wrong way round Mg-rich. The gradual oxidation of the aqueous Fe2+ species eventually leads to precipitation of magnetite, which is also metastable because the aqueous solution is maximally supersaturated with respect to hematite. In most computer implementations, the aqueous activities of Fe2+ and Fe3+ are adjusted to redox equilibrium and the supersaturation with respect to Fe-bearing solids is determined from these activities. Unfortunately, this approach does not provide supersaturation values in ‘real time’. Reactions between aqueous species are actually not instantaneous and can be slower than the dissolution/precipitation of related solids. Considering this fact is essential in modeling the artificial carbonation of ferromagnesian minerals for CO2 sequestration purposes.

1

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1. INTRODUCTION

Solid solutions between carbonate minerals have attracted attention from the scientific community since the second half of the twentieth century.1-5 The need for sustainable remediation procedures (immobilization of harmful ions, CO2 mineralization, etc.) has progressively put carbonate minerals and their solid solutions in the scientific spotlight.6-9 For example, the injection into basaltic rocks of CO2 bubbles within down-flowing water seems to accelerate metal release from basalt, with the subsequent crystallization of carbonate minerals.10 The crystallization behavior in the (Mg,Fe)CO3–H2O system deserves a deep study, since this solid solution is expected to be one of the main phases formed during carbonation processes of mafic silicate minerals. In most mafic minerals (olivine, pyroxene, amphibole, etc.), the substitution of Mg for Fe(II) is complete (absence of miscibility gaps for intermediate compositions in the series) and this seems also to be the case in the MgCO3-FeCO3 series.11-15 In fact, a key factor in determining the miscibility between substituting ions in a mineral structure is the difference between their effective cationic radii4, which for (Mg,Fe)CO3 solid solution is ∆( ) ≈ 0.06 Å.16 This value is significantly

smaller than those of systems with limited miscibility, such as (Ca,Fe)CO3, where ∆( ) ≈

0.22 Å or (Ca,Mg)CO3 with ∆( ) ≈ 0.28 Å. The term breunnerite is frequently used to

refer to a variety of ferroan-magnesite with an Mg/Fe ratio from 90/10 to 70/30, but this mineral name is not included in the IMA (Int. Min. Assoc.) list, due to the complete miscibility between siderite and magnesite. While the thermodynamics of the (Mg,Fe)CO3–H2O system is well established, there are few data on the effective crystallization mechanisms in this system under different conditions of temperature and pressure. Here, in order to fill this lack, we study the crystallization-dissolution behavior of different members of the (Mg,Fe)CO3 solid solution during hydrothermal cycles performed at 150 and 180 ºC (5-9 bars). The reacting solutions were supersaturated with respect to magnesite (MgCO3), siderite (FeCO3), and all the members of the (Mg,Fe)CO3 solid solution. Previous experiments have been usually performed at temperatures in the range 350-500 ºC and at 2-3 kbars1,3,12,14 but the aim of this paper is studying the formation and evolution of (Mg,Fe)CO3 at lower temperatures and pressures, which are more relevant from the point of view of CO2 sequestration. We have performed experiments at high supersaturation with the aim of obtaining precipitates with an initial Mg/Fe ratio close to that in the parent solution and studying their later 2


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evolution under epithermal conditions. We are particularly interested in ascertaining the influence of Fe(II)-Fe(III) processes on the crystallization behavior of ferromagnesian carbonates, namely: (i) To which extent the precipitation of Fe(III)-solids is an iron sink in aqueous systems without contact with the atmosphere? (ii) What mechanisms determine the growth kinetics of (Mg,Fe)CO3? Answering these questions is an important challenge in the task of modelling the artificial carbonation of ferromagnesian minerals in injection sites. Finally, we discuss briefly the implications of including Ca2+ (in addition to Mg2+ and Fe2+) in the scenario.

2. THERMODYNAMIC BACKGROUND AND EQUILIBRIUM CALCULATIONS 2.1. Equilibrium Mg2+/Fe2+ partitioning Chai and Navrotsky14 proposed a non-ideal, “strictly” regular model2 for the substitution of Mg2+ for Fe2+ and determined a positive enthalpy of mixing for the (Mg,Fe)CO3 solid solution. In such a model, the enthalpy of mixing (∆HM) is a symmetric function of the solid composition that has a maximum for Mg0.5Fe0.5CO3. Experimental data of the mixing properties of non-ideal solid solutions are usually fit to a Guggenheim expansion series,13 which in the simplest case of a regular solid solution becomes: ∆ = !"

(1)

where R and T are the gas constant and the absolute temperature, XMgs and XSid = 1-XMgs the mole fractions of the magnesite (Mgs) and siderite (Sid) components in the solid solution, and a0 is a dimensionless Guggenheim parameter. Chai and Navrotsky14 give a value of 4.44 kJ/mol for the functionally equivalent Margules parameter (W = a0RT). In the case of a ‘strictly’ regular solid solution, a0RT does not depend on temperature and pressure (see errata correction of reference17), which implies a null excess volume (∆VE = 0) and entropy (∆SE = 0) of mixing. Therefore, the excess free energy of mixing ∆GE = ∆HM and is given by Eq. (1). Finally, the solid-phase activity coefficients (γ) are functions of ∆GE that relate the activity of the solid solution components to their molar fraction.16 In the case of a regular solid solution: & ln % =

(2)

3



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and & ln % =

(3)

When we know the mixing properties of the solid solution, calculating the equilibrium (eq) distribution of the substituting ions between the aqueous (AQ) and the solid solution (SS) is straightforward:19,20 '()*-

+, (,.)

=

/)*0 ⁄/123

5 6⁄5 6(7)

8

:

= 8 123 ∙ : 123 )*0

)*0

(4)

where the terms in curly brackets represent the activities of the corresponding free (uncomplexed) aqueous ions. The quotient KSid/KMgs between the solubility products of the two end-members depends on temperature but is independent on composition. Differently, γSid/γMgs accounts for the non-ideality of the solid solution and depends on the solid composition. Figure 1a displays the solubility products of magnesite and siderite as a function of temperature, calculated using the analytical expressions determined by Bénézeth and co-workers.21,22 Figure 1b shows the variation of D{Mg/Fe}(eq) as a function of both temperature and composition. Note that D{Mg/Fe}(eq) is significantly smaller than 1, which means a strong tendency of Fe2+ to incorporate into the solid phase, D{Fe/Mg}(eq)= 1/D{Mg/Fe}(eq) >> 1, because the lower solubility of siderite. It is also worth noting the high variability of D{Mg/Fe}(eq) with the solid composition for a given temperature. As can be seen, the solubility of both end-members decreases dramatically at increasing temperature but the decreasing rate is higher in the case of magnesite. As a result, the quotient KSid/KMgs increases from 0.008 at 25 ºC to 0.01 at 180 ºC and D{Mg/Fe}(eq) increases as well, which is in agreement with the crystallization behavior observed in pioneer experimental works.12

Figure 1. (a) Variation of the solubility products (logK) of magnesite and siderite with temperature: KMgs decreases from 10-7.8 (at 25 ºC) to 10-11.4 (at 180 ºC). KSid decreases as well in this range of 4


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temperature (from 10-10.9 to 10-13.4) but at a lower rate. (b) Variation of D{Mg/Fe}(eq) as a function of temperature and solid composition (XSid). 2.2. Equilibrium solubility diagrams

Equilibrium in AQ-SS systems is frequently described using the so-called “Total Activity Product” defined by Lippmann.23 For the (Mg,Fe)CO3-H2O system: ΣΠ = (5=>&? 6 + 5AB &? 6)5CDE& 6

(5)

The curly brackets stand for the aqueous activities of the corresponding ions. The “Total Solubility Product” (ΣΠ7 ) represents the value of ΣΠ at thermodynamic equilibrium, which expressed as a function of the solid composition becomes: ΣΠ7 = F + F = % F + % F

(6)

where = % is the activity of the siderite component in the solid phase. An analogous expression defines aMgs. When expressed as a function of the aqueous composition, the total solubility product is given by: /+,,J.

ΣΠ7 = 1LH8

123 :123

+8

/)*,J.

)*0 :)*0

K

(7)

where ,7 = 5AB &? 6⁄(5=>&? 6 + 5AB &? 6) is the activity fraction of Fe2+ in the aqueous solution and XMg,aq = 1- XFe,aq. Equations 6 and 7 were denominated solidus and solutus by Lippmann21 and their graphical representation yields a diagram in which solidus and solutus are represented on the ordinate against two superimposed, XSid and XFe,aq, scales on the abscissa. A detailed derivation of Eqns. 6 and 7 can be found in the work by Glynn and Reardon.24 Horizontal tie-lines can be drawn between the solidus and solutus curves to obtain the aqueous and solid solution compositions at thermodynamic equilibrium. Figure 2 shows the Lippmann diagram calculated in this work for the (Mg,Fe)CO3–H2O system at 150 ºC. As can be seen, due to the high preferential partitioning of iron towards the solid phase, intermediate members (XSid < 0.9) of the solid solution are at equilibrium with Fe2+-poor (XFe,aq < 0.05) aqueous solutions. In practice, kinetic effects can lead to significant deviations of the effective distribution coefficients from these equilibrium values, but a precise knowledge of the equilibrium 5



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thermodynamics is the key to account for non-equilibrium phenomena. The deviation typically results in a reduction in the degree of preferential partitioning, the solid solution becoming richer in the more soluble component20,25-27 and the distribution coefficient approaching unity. In the long term, the system will tend to reach equilibrium through a dissolution-recrystallization process, but kinetic and mechanistic factors can inhibit this process by increasing the lifetime of transitory phases. This is the case of numerous AQ-SS systems in which the extremely slow recrystallization kinetics can make metastable compositions to persist for a long time.

Figure 2. Lippmann diagram for the (Mg,Fe)CO3–H2O system at 150 ºC. The range of aqueous activity fractions of iron that are at thermodynamic equilibrium with solids of composition XSid < 0.9 is very narrow (XFe,aq < 0.05) as underlined by the ellipse on the upper left side of the graph.

3. MATERIALS AND METHODS 3.1. Precipitation experiments Precipitation experiments were performed by mixing solutions of MgCl2, FeCl2, and Na2CO3 following the relation: M=>NO + MABNO = "PC = 0.1 QRS/U>V

(8)

where the superscript i stands for “initial” and the subscript T refers to the total analytical concentration of the corresponding element (into brackets). Analytical degree chemicals (SigmaAldrich) and ultrapure (MilliQ; 18.2 MΩ·cm) water were used in all the cases. The initial 6


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concentrations of the reactants are shown in Table 1, where the experiment codes (A, B, C, D, E, F, and G) correspond to the 7 different M=>NO ⁄MABNO ratios in the initial solution. The reactants were dissolved in 60 g of water and the resulting solutions were added to a polytetrafluoroethylene (PTFE) vessel (Berghof) in the order Mg, Fe, C. Finally, we added 48 g of water to reach a total amount of 228 g. The total inorganic carbon concentration (TICi) in all the different parent solutions was 0.1 molal and the filling factor of the vessel was 75% in volume. In order to minimize the oxidation of Fe2+ to Fe3+, the ultrapure water was pre-treated to reduce the amount of dissolved oxygen by heating up to the boiling point for at least half an hour while N2 was bubbling inside.26 Then, the solution was rapidly cooled with iced water and stored (during less than 2 hours) in a glove box (Plas-Labs Inc.) under nitrogen atmosphere. The PTFE vessel was then inserted into a stainless steel high-pressure reactor (Berghof, BR-300) equipped with a temperature controller (Berghof, BDL-3000), an analogue manometer and a magnetic stirrer (impeller diameter Øimp = 2 cm). Before heating, N2 was flowed through the reactor for 15 minutes to keep a nitrogen atmosphere. The reactor was heated and kept at the target temperature under continuous stirring (750 rpm) for defined periods of time. Finally, after a cooling period without stirring, the solutions were extracted, filtered (Ø = 0.45 µm, Pall Corporation), and analyzed for magnesium and iron by ICPMS (Agilent Technologies, HP 7500c). A set of standards, prepared using certified standard solutions (Perkin Elmer, Norwalk, CT) was run before and periodically during each sampling event.

Sc, Y, Rh, In, and Th were used as internal standards. The analytical precision was better than 4% (RSD) in all cases. Alkalinity (±1% RSD) and pH (±0.05) were measured with a CRISON Compact Titrator combined with an Orion electrode, and calibrated with buffer solutions of pH 4.00, 7.00 and 9.21 (Crison). Each experimental run was conducted in duplicate (some in quadruplicate), thus, the values reported in Table 1 are the average of parallel experimental runs and the uncertainties (shown on top of the corresponding columns) involve the whole experiment. Note the high uncertainty of the final iron concentrations in the aqueous solution. A detailed discussion about these values is given in section 4.5. Geochemical modeling of the aqueous solution was carried out from the analytical data using the geochemical code PHREEQC.29 It is worth noting that the experimental set up used in this work is a kind of ‘black-box’ in which we know the initial and final conditions (reactant concentrations, pH, etc.) but we ignore the evolution of fluid and solid phases all along the hydrothermal cycle. Most hydrothermal experiments are not really isothermal since include heating and cooling stages. 7



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Here, every cycle involves: (i) A heating period of about 90 minutes to reach the target temperature (150 or 180 ºC and 5-9 bars), (ii) a period (1 to 30 days) at constant temperature, and (iii) a period of about 3.5 hours in which the reactor is left to cool and depressurize to 25 ºC and 1 bar. Given that the solubility increases with decreasing temperatures, the precipitate could start to dissolve during the cooling period. However, the kinetics do not seem to be very favorable for that, especially because the cooling period was performed without stirring.

3.2. Characterization of the solid phases

The precipitates were dried at 35°C and characterized by powder X-ray diffraction (XRD; PANalytical X’Pert Pro, Cu-Kα) using Silicon (NIST 640b) as internal standard to correct the peak positions. The samples were analyzed in the range 5-80 º2ϴ, with a step size of 0.013 º2ϴ and a measuring time per step of 39.27 seconds. The diffractograms were studied using X’Pert HighScore Plus 2.2.4 (Panalytical). Accurate peak profile analysis was performed after Kα2 stripping, according to Rachinger method.30 The mole fraction (XSid) of siderite in the solid solution was determined using the d-spacing of the 104 reflection. This parameter has been proved to vary linearly with composition from pure siderite to pure magnesite.14,31 The compositional homogeneity was checked by measuring the full width at half maximum (FWHM) of the main reflections. In the case of the pure end-members, the crystallite size (the size of a coherently diffracting domain) was also calculated from the FWHM values using the Debye-Scherrer tool of X’Pert HighScore Plus. Scanning electron microscope (SEM-EDX, JEOL-6610LV) and electron microprobe (EPMA, CAMEBAX SX-100) analyses were used to characterize the morphology, homogeneity and composition of several individuals in each sample and to obtain composition profiles on polished sections of crystals exhibiting compositional zoning. Moreover, X-ray photoelectron

spectroscopy (XPS) analyses were carried out to evaluate the Fe2+/Fe3+ proportion in some representative samples. With this aim, the precipitates were crushed to fine powder, homogenized and pressed into a tablet 1 cm2 in size and up to 4 mm thick. Thus, the values obtained were expected to be representative of the average Fe2+/Fe3+ ratio in the precipitate. To avoid reduction of Fe2O3 to Fe3O4 by Ar+, the samples were analyzed without previous treatment by argon bombardment. The analyses were carried out with a SPECS Phoibos 100 MCD5 instrument, using a beam of non-monochromatic magnesium X-rays (1253.6 8


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eV photon energy), with an electric potential of 13 kV and a power of 250 W. Due to the low conductivity of the samples the charge effect was compensated using a low energy electron source. The complete scanning was carried out using a 1 eV step energy, and a pass energy of 90 eV in a single scan. The high-resolution spectra for each element were done with a pass of 0.1 eV and with an energy pass of 30 eV, averaging 5 or 10 times as seen necessary. The background pressure in the analysis chamber was kept below 5×10−9 mbar. The spectra were corrected with the C1s peak fixed at 284.6 eV as a reference for charge correction. The chemical state of iron was determined from the binding energies (BE) of Fe2p3/2. With this aim, peak fitting and deconvolution procedures were accomplished using CasaXPs processing software.

4. RESULTS AND DISCUSSION 4.1. Final composition of the aqueous solution in the precipitation experiments Table 1 displays the experimental conditions and the initial (i) and final (f) concentrations of the reactants, [Mg], [Fe], etc., in the aqueous solution. Every row in Table 1 displays the mean values of replicate experiments. The final, total inorganic carbon concentration (TIC) was calculated from pH and alkalinity at 25 ºC using PHREEQC.29 As discussed in section 2, when crystallization occurs near equilibrium, the precipitation of intermediate (Mg,Fe)CO3 members requires a low proportion of Fe2+ to be present in the aqueous solution (Figure 2). However, our experiments were performed at high supersaturation with the aim of obtaining precipitates with an initial Mg2+/Fe2+ ratio close to M=>NO ⁄MABNO and then checking the effect of the heating time on the subsequent evolution of the system towards equilibrium, emulating processes that can be relevant from the point of view of CO2 sequestration. An amorphous, precursor precipitate (see inset in Figure 3) forms instantaneously, as soon as the reactant solutions are mixed in the PTFE vessel, before starting the heating cycle. The solution pH changes quickly during this process so that initial pH compiled in Table 1 has been calculated from the concentration values of the reactant solutions. The saturation index (SI) of the hypothetical (homogeneously mixed but unreacted) initial solution with respect to magnesite, siderite, and the solid solution has been calculated by applying the speciation

9



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code PHREEQC to the data displayed in Table 1. In the case of magnesite, the saturation index is given by: WP = log(PZ[⁄F ),

(9)

where PZ[ = 5=>&? 65CDE& 6 is the ionic activity product in the initial, supersaturated solution. An analogous expression defines SISid. The concept of supersaturation with respect to intermediate members of the solid solution is complex and beyond the scope of this work. A detailed discussion can be seen in a recent review.7 For the present purposes, we can use the concept of stoichiometric saturation index (SIst), given by the function: 5 6(^_`) 5 6` 5ab _ 6 ` (^_`) 123 :123 c) (8)*0 :)*0 (dc))

WP\ (]) = SR> (8

(10)

where ] = = 1 − . Equation 10 represents an useful estimation of the supersaturation in AQ-SS with some features that need to be considered:7 (1) Strictly, Eqn. 10 is valid for congruent dissolution or growth in which the transfer of matter between solid and fluid occurs in the same stoichiometric proportion as in the solid formula, i.e. FexMg(1-x)CO3. (2) For a given aqueous solution, the function in Eqn. 10 represents the departure from “stoichiometric equilibrium” for all solid-solution members.32 (3) The curve has a peak at the solid composition (Xmax) with respect to which the aqueous solution is maximally supersaturated. (4) Solid solutions with compositions for which SIst > 0 can grow congruently. (5) Solid solutions with compositions for which SIst < 0 can dissolve congruently. (6) When a solid of composition Xmax precipitates, the effective distribution coefficient equals the equilibrium value. Table 2 shows the saturation indices (at 25 ºC) with respect to magnesite, siderite, and (Mg,Fe)CO3 fc ) of the solid with in the hypothetical (unreacted), initial solution, as well as the composition (

respect to which the initial aqueous solution is maximally supersaturated. The initial solutions in the same series (A, B, C, D, E, F, or G) are identical. All of them plot above the ‘solidus’ and are supersaturated for any composition in the series. In all cases, the highest supersaturation correspond to a Fe-rich (XSid > 0.97) member of the (Mg,Fe)CO3 solid solution.

Table 1. Initial and final concentrations (mmol/kgw). In all cases TICi = 100 mmol/kgw.

10


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Exp.

MghNji

MklNji

pHi

T (ºC)

Time (min)

MghNmi ±10%

MklNmi ±35%

TICf ±20%

pHf ±3%

XSid-1 ±10%

XSid-2 ±20%

nopjqr nopjqs

A1

90

10

10.5

180

1440

1.678

0.005

4.236

6.72

0.09

-

-

B1

80

20

10.4

180

1440

1.822

0.010

4.708

6.72

0.20

0.09

1.94

B2

80

20

10.4

180

20160

2.269

0.002

3.668

6.61

0.27

0.06

0.69

B3

80

20

10.3

180

43200

3.402

0.001

5.494

6.68

0.14

-

-

C1

60

40

10.3

180

1440

3.181

0.218

5.078

6.77

0.42

0.19

2.49

C2

60

40

10.3

180

10080

2.199

0.001

5.488

6.78

0.44

0.04

1.93

C3

60

40

10.3

150

10080

1.805

0.125

3.133

6.66

0.49

0.15

2.05

D1

50

50

10.2

180

1440

2.156

0.181

5.418

6.59

0.54

0.26

3.46

D2

50

50

10.2

120

10080

2.694

0.010

6.769

6.49

0.50

0.34

1.14

D3

50

50

10.2

150

7200

1.671

0.112

3.212

6.50

0.53

0.30

3.19

D4

50

50

10.2

150

10080

2.267

0.113

4.129

6.59

0.53

0.27

3.36

D5

50

50

10.2

150

14400

2.591

0.033

4.513

6.63

0.49

0.26

4.36

D6

50

50

10.2

150

30000

3.421

0.024

5.422

6.55

0.53

0.17

3.37

E1

40

60

10.2

180

10080

3.156

0.001

6.468

6.40

0.58

0.44

1.78

F1

20

80

10

180

20160

2.159

0.377

5.427

6.27

0.74

0.50

10.2

G1

10

90

9.93

180

1440

0.620

0.402

5.255

5.98

0.90

-

-

G2

10

90

9.93

180

10080

0.948

0.116

5.492

5.98

0.88

-

-

Table 2. Initial saturation indices with respect to the magnesite, siderite, and (Mg,Fe)CO3 (mmol/kgw)

MklNji

(ttuv/wxy)

pzjgh{

pzjpjq

pzjpp

on|} pjq

A

90

10

3.59

4.44

4.46

0.978

B

80

20

3.50

4.77

4.77

0.991

C

60

40

3.34

5.08

5.08

0.996

D

50

50

3.24

5.18

5.18

0.999

E

40

60

3.20

5.26

5.26

>0.999

F

20

80

2.74

5.38

5.38

> 0.999

G

10

90

2.38

5.42

5.42

>0.999

Sample

MghNji

4.2. Characterization of the solid phases Figure 3 shows the diffraction patterns of some typical precipitates obtained in these experiments. Peak broadening is significantly higher for the siderite end-member than for magnesite, which is in agreement with the increasingly flaky aspect of the crystals in the SEM 11



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Page 12 of 30

images. For magnesite, the crystallite size ⊥ {104} was ~250 nm, whereas for siderite was ~150 nm. The series is known to satisfy the “Vegard Rule” and so a linear interpolation allows estimating the composition (XSid) from the d-spacing of specific reflections.14, 31 The most suitable is 104 (the strongest one) that has a shifting ∆d104 = 0.053 Å between the end-members.

Figure 3. Typical diffraction patterns of some samples synthesized in this work. We have removed the background to better show the shifting of the peaks with composition. Reflections occur at increasing 2θ (decreasing spacing) as the magnesium content increases. The inset on the top corresponds to the precursor, amorphous phase. SEM images of the corresponding precipitates are also shown.

A problem arises because all the diffractograms exhibit double reflections, being the doubling of 104 particularly clear due to its intensity. The presence of wide ‘double-peak’ reflections indicates compositional heterogeneity. Figure 4 illustrates this effect in the case of the sample C3b. Note that the code C3 in Table 1 involves four experiments labelled C3a, C3b, C3c, and C3d. The phenomenon is completely repeatable, as can be seen in Fig. 4d, which shows the doubling of 104 in these four samples. The compositions corresponding to both peaks are shown in the inserted table, the mean values being XSid-1 = 0.49 ± 0.03 and XSid-2 = 0.15 ± 0.04. Both values differ 12


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significantly from the composition obtained by semi-quantitative EDX-microanalyses as can be seen in Fig. 4e, which shows the EDX-composition of 5 different crystals of the C3b sample. However, the mean EDX-value, XSid = 0.38 ± 0.06, falls within the range from XSid-1 to XSid-2 and could be considered representative of the whole composition.

Figure 4. Compositional heterogeneity and ‘double-peak’ XRD reflections in sample C3b. (a) Scan data-points (step size 0.013 º2ϴ) after Kα2 stripping. (b) Experimental profile (red line) and pseudoVoight fitting function convoluting both peaks (blue line). The peak broadening value (FWHM) is also given. (c) Experimental profile (red line) and pseudo-Voight fitting function considering two separate peaks (blue line). Values of composition corresponding to both peaks (XSid-1 and XSid-2) are given. (d) Peak doubling in four analogous experiments. The compositions corresponding to both peaks are shown in the inserted table. (e) Mean composition obtained from 5 semi-quantitative EDX analyses (marked on the SEM image of the precipitate).

4.3. Why does peak doubling occur?

In absence of a thermodynamic miscibility gap, the development of a bimodal crystallization behavior can only be attributed to the ‘crystallization history’ in the system. There are two processes that can produce this effect: Development of a sharp compositional zoning within each single crystal or coexistence of two precipitates of different composition. Both effects are favored in AQ-SS systems in which the solubility products of the end-members differ by 2 or 3 orders of 13



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magnitude,33 which is the case here. The question is when compositional switching occurs in our ‘black box’, hydrothermal experiments. Figure 5a gives some keys to reveal such a riddle. In this figure the initial mole fraction of iron in the aqueous solution has been plotted against the final composition of the solids, either XSid-1 (the composition calculated by the position of the main 104 reflection in Fig. 4) or XSid-2 (the composition corresponding to the secondary 104 reflection). After an inspection of Fig. 5a, an immediate conclusion arises: All the black dots (XSid-1 data-points) lie close to the diagonal from (0,0) to (1,1), while the red stars (XSid-2 data-points) plot below this diagonal. This means that the initial composition of the aqueous solution directly determines the composition of the XSid-1 solids, which in turn has two significant implications: First, at the high initial supersaturation (see Table 2), the substituting ions (Fe2+ and Mg2+) are incorporated in the precursor, amorphous precipitate in a stoichiometric proportion close to that of the aqueous phase. Second, the XSid-1 data-points correspond to crystalline precipitates formed in the beginning of the heating cycle from the remaining fluid and the solvent-mediated recycling of the amorphous precursor. Note that during the initial 90 minutes, the temperature is continuously increasing. Therefore, the supersaturation level remains high since the solubility of both end-members decreases at increasing temperature (see Fig. 1a). At high supersaturation and precipitation rate, there is less opportunity for thermodynamic selectivity effects to be exerted (particularly when an unstable, much more soluble precursor occurs) and the effective distribution coefficient deviates towards unity. A comparison between the fc column in Table 2) and compositions predicted from maximum supersaturation criteria (see

the effective compositions (XSid-1 column in Table 1) makes clear this effect. The fact that the iron content of the precipitate (XSid-2 column in Table 1) becomes subsequently lower is not a simple kinetic effect. Actually, since magnesium and iron tend initially to incorporate equally into the solid phase, their proportion in the aqueous solution keeps approximately constant. Afterwards, as the supersaturation decreases, the thermodynamic driving should prevail, i.e. the solid solution should become increasingly rich in iron. However, the outcome is just the opposite! The sample inhomogeneity can be estimated by plotting FWHM [º2ϴ] against the composition (Figure 5b). In this graph, the composition corresponding to each data-point was calculated from the maximum of a convoluted peak that fits the overall perimeter as in Fig. 4b. The position of this maximum represents an intermediate composition between XSid-1 and XSid-2 (see Fig. 4c). The marked compositional inhomogeneity of the intermediate precipitates is evident in Figure 5b. The high FWHM values observed for intermediate compositions, together with the scattered 14


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arrangement of the data-points, indicates the presence of two populations in these samples. The molar ratio between these two populations can be estimated by integrating the areas of each of the two 104 reflections and it is displayed in Table 1 (last column). As we will see, the total amount (in moles) of both populations can be calculated by mass balance of carbonates. However, we have to consider previously that a fraction of dissolved iron and minor amounts of magnesium incorporate into secondary, non-carbonate phases.

Figure 5. (a) Comparison between the initial mole fraction of iron in the aqueous solution and the final composition of the solid solution (Mg,Fe)CO3 determined from the position of the 104 reflection. (b) FWHM versus the mole fraction of Fe2+ in the solid solution. Note that the intermediate members exhibit a stronger inhomogeneity than Fe-rich and Mg-rich members. To finish this section, we should remark that absence or insignificance of the second 104 reflection occurs when the concentration of either magnesium or iron in the initial aqueous solution is relatively high. Such is the case of the experiments in the series A, B, and G. In the first case (A and B), the solid solution chemistry is initially controlled by magnesium and there is no switching to a second crystallization period because, in the long term, the available iron incorporates into secondary phases.

In the second case (G), the solid solution chemistry is initially controlled by

iron and when the supersaturation decreases, the strong tendency of iron to partition into the solid solution keeps the precipitate composition within the iron-rich (XSid > 0.9) range. It is also worth noting that, although in most cases peak doubling is due to the presence of two crystal populations, concentric compositional zoning in single crystals has occasionally been observed. Figure 6a shows the backscattered electron image of a polished section of a large solid solution crystal. A compositional profile obtained by electron microprobe is also shown. As can be observed, the core is richer in iron whereas the rim is rich in magnesium. This sample correspond to one of the 15



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Page 16 of 30

replicate experiments labelled D5 in Table 1. Interestingly, the surface of some crystals is covered by a third population (Fig. 6b) consisting of small crystals with a typical rhombohedral shape (2-10 µm) that seems to be formed in the late stages of the experiments. These crystallites exhibit also concentric zoning, but their core is richer in magnesium.

Figure 6. (a) Backscattered electron image of a polished section of a large solid solution crystal. The crystal exhibit a sharp concentric zoning with compositions analogous to those detected by Xray diffraction. (b) A third generation of solid solution crystals has occasionally been observed. These crystals exhibit compositional zoning as well, but in this case the core is richer in magnesium. Sample taken from one of the replicate experiments labelled D5.

4.4. Secondary phases Whereas the (Mg,Fe)CO3 solid solution is the prevailing solid phase, there are some important details in the diffractograms that need to be considered. The first one is related to the background level. In the sequence of XRD patterns in Figure 3, the background was removed in order to better display the sifting of the peak positions. However, as shown in Figure 7a, all the diffractograms exhibit a significant background, the higher the iron content, the higher the background level. Such an increasing background indicates the presence of iron-rich low-crystallinity phases, likely Fe(III) oxy-hydroxides. Moreover, the presence of a crystalline phase other than (Mg,Fe)CO3 has been observed in a fraction of the precipitates that strongly adheres to the magnetic stirrer. Figure 7b shows a diffractogram of this phase, identified as magnetite-magnesioferrite (Mag-Mfr). There is a solid

solution

series

between

magnetite,

[Fe3+]t[Fe3+Fe2+]oO4,

and

magnesioferrite, 16


Page 17 of 30

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[Fe3+]t[MgFe3+]oO4, with a A~3Q spinel-type structure and extremely close unit-cell parameters. Both end-members are strongly magnetic and the stirrer works as a magnetic separator. The superscripts t and o stand for tetrahedral and octahedral coordination. These occupation schemes represent the stable arrangement at low temperature. At high temperature cations exchange between tetrahedral and octahedral sites, thereby producing a non-convergent, disordered structure with the same symmetry.34,35

Figure 7. (a) Powder XRD diagrams illustrating the influence of the increasing iron content (from A1 to G1) on the background level. (b) Diffractogram of the magnetically-separated Mfr-Mag fraction. Note that the reflection 111 could be overlapped with the main reflection of PTFE.36 The typical double-peak of the (Mg,Fe)CO3 precipitate is also present. 4.5. Fe(II)-Fe(III) behavior: An inextricable mess?

As previously discussed, the equilibrium thermodynamics of the (Mg,Fe)CO3-H2O system indicates that there is a preferential tendency of Fe2+ to incorporate into the solid phase (see Fig. 2). However, at the high initial supersaturation level, the substituting ions tend to incorporate into both the amorphous precursor and the initial carbonate solids in a stoichiometric proportion close to that in the aqueous phase (see Table 1). The high Mg-content of the initial solids is mostly due to ‘unselective’ substitution at high precipitation rates. Moreover, low crystallinity phases obtained at high supersaturation exhibit anomalous inter-planar distances that also favor the indiscriminate incorporation of both ions, especially when they have the same charge and similar atomic radii. Obviously, both the amorphous precursor and the early crystalline phases are not at equilibrium with the remaining solution and the system will tend to undergo a dissolution-recrystallization process to reach equilibrium. This solvent-mediated transformation will be favored by the low crystallinity of the initial precipitates as reported for other carbonate systems.37-39 In the case of Mg17



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Page 18 of 30

bearing carbonates the problem is particularly complex due to the number of metastable hydroxycarbonate-hydrates that could potentially form during the early stages of the heating period.40,41

The presence of Mfr-Mag solids suggests that the second (Mg,Fe)CO3 precipitation event is driven not only by the lower solubility of the Fe-rich members of the solid solution but also by the oxidation of a fraction of Fe(II) to Fe(III). We can expect two main reactions to compete for the ferrous Fe2+ ions, namely: &? & &? &? &? Mg (7) + Fe&? (7) + COE(7) → My − (1 − x)NMg (7) + ( − ])Fe(7) + Mgdc Fec COE()

(11)

and &? E? \ E? &? Fe&? (7) + O&( ) + 6OH(7) → MFe N MFe Fe N O() + 3H& O() + ( − 3)Fe(7) .

(12)

Note that in Eqn. (12) we do not consider cation exchange between tetrahedral and octahedral sites in the Mag-Mfr solids, given the moderate temperature of the experiments.35 Moreover, as we discuss later, mass balance calculations indicate that the incorporation of magnesium into Mag-Mfr was negligible in all the experiments. Equations (11) and (12) only involve the solid phases identified by XRD. Formation of other Fe(III)-bearing phases of very low solubility, particularly hematite (Hem), can be expected to occur. However, when present in minor amounts, these phases are difficult to detect by XRD because of its low crystallinity. Iron oxidation in water is an extensively studied process42-44 of which the kinetics can be estimated with a certain level of confidence. Kinetic studies indicate that conversion of Fe(II) into Fe(III) depends dramatically on the solution pH, which determines the distribution of the Fe(II)bearing aqueous species. In the pH range (6-7; see Table 1) of our experiments, Fe(OH)20 can be expected to be the prevailing species in absence of inorganic carbon and other complexing and ion pairing agents. Fe(OH)20 determines the overall rate because it is very much readily oxidized than FeOH+ and Fe2+.44 However, as a good number of authors44,45 have already pointed out, the presence of Cl- and CO32- retards significantly the oxidation rate because these ions form Fe(II)-bearing complexes that are slowly oxidized in comparison to Fe(OH)20. In our case, HCO32- and CO32- ions are present at high concentrations and the solubility of ferrous iron tend to be controlled by (Mg,Fe)CO3. Moreover, there is no external supply of oxygen and the concentration of ironchlorine complexes is leading, due to the high concentration of Cl- in the parent solutions (Table 1). Therefore, the amorphous precursor can be expected to be free of Fe(III) because its precipitation 18


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rate is orders of magnitude higher than the expected oxidation rate. In contrast, the solventmediated transformation of the initial precipitate to form the second population of (Mg,Fe)CO3 occurs at a slower rate and coexists with the precipitation of Fe(III)-bearing phases like magnetite and/or hematite.

The previous discussion explains in a qualitative way the crystallization behavior in the present experiments and, particularly the influence of the iron oxidation in the process. The deficiency of Fe(II), produced by oxidation to Fe(III), results in a growth stage in which the carbonate solid solution (Mg,Fe)CO3 becomes richer in magnesium despite some apparently contradictory facts: (i) According to the (Mg,Fe)CO3-H2O thermodynamics, iron would tend to incorporate preferentially into the solid solution, particularly in near-to-equilibrium conditions. (ii)

Once the target

temperature is attained, the system approaches equilibrium and the thermodynamic selectivity effect will tend to promote the preferential incorporation of Fe towards the solid. (iii) Despite the high supersaturation, in the first precipitation stage there is a slight preferential partitioning of Fe towards the (Mg,Fe)CO3 solid solution (black dots plot close to diagonal in Figure 6 but their global tendency is slightly Fe-rich). Oxidation of Fe(II) and subsequent precipitation of Fe(III)-bearing phases is a sink for Fe(II) and explains the decrease of the Fe/Mg ratio in the carbonate solid solution formed in the second stage. The experiments D3-D6 were all performed at 150 ºC from the same parent solution and are useful to illustrate the reaction progress as a function of time. Figure 8 shows the composition evolution of both the aqueous and the solid solution, plotted on the corresponding Lippmann diagram. There is an initial, dramatic decrease in the solution concentration but after 5 days of reaction, ΣΠ slightly increases and the aqueous solution becomes richer in Mg2+. The aqueous activity fraction XFe,aq decreases accordingly during this period to reach a value of 0.0024 after 21 days of reaction. The compositional evolution of the (Mg,Fe)CO3 solids follows the same trend. In the early stages the Fe/Mg ratio in the solid is close to the Fe/Mg ratio in the parent solution (XSid-1 ~0.5), but the subsequent process leads to formation of a second precipitate poorer in iron (XSid-2 < 0.3). Such an evolution occurs despite the aqueous solution is maximally supersaturated with respect fc to Fe-rich solid solution members: The horizontal lines in Fig. 8 connect solidus and solutus

XFe,aq compositions at thermodynamic equilibrium. The vertical lines connect saturated and fc supersaturated aqueous solutions with the same value of XFe,aq. In each case, is the solid

composition (0.96, 0.94, 0.68, and 0.31) for which the stoichiometric supersaturation function has a 19



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Page 20 of 30

maximum (for a detailed explanation see reference7) at the corresponding reaction time (5, 7, 10, fc and 21 days). Note that in all cases & , i.e. the experimental values are richer in

magnesium.

Figure 8. Evolution of the total ionic product with the reaction time (0, 5, 7, 10, and 21 days). Reaction temperature = 150 ºC, Fe/Mg ratio = 1, experiments D3-D6. The L-lines (green, purple, blue, and orange) connect the composition of the experimental solution with the composition of the solid with respect to which the aqueous solution is maximally supersaturated. The diamonds at the extremes of the horizontal segments correspond to the two experimental values of the solid composition. Figure 9 shows the evolution of the aqueous concentrations of Fe, Mg, and TIC and allows qualitatively explaining the crystallization sequence. After the first strong fall, there is an increase of concentration that can be explained by dissolution of the initial, metastable (Mg,Fe)CO3 solids coupled with the formation of (Mg,Fe)CO3 richer in magnesium and secondary non-carbonate phases (Mag-Mfr and amorphous oxides). Further compositional arrangements keep TIC and Mg concentrations slightly increasing until the end of the experiment. Figure 9b displays the driving force for this dissolution-recrystallization process: After ~3 days, the aqueous solution becomes

undersaturated with respect to intermediate compositions (XSid ~5) and supersaturated with respect

to Mg-rich compositions of the solid solution. In contrast, the concentration of iron continues decreasing after the initial fall, due to Fe(II)-Fe(III) oxidation and subsequent precipitation of magnetite and other oxides. As shown in Figure 9c, the saturation index for the oxides is significantly higher than for the carbonates in all the experiments.

20


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Figure 9. (a) Evolution of the aqueous concentration of Mg, Fe, and TIC. Reaction temperature = 150 ºC, initial Fe/Mg ratio = 1, experiments D3-D6. (b) Evolution of the stoichiometric saturation index of two different compositions of the (Mg,Fe)CO3 solid solution. (c) Evolution of the saturation index of magnetite and hematite. The previous sequence is absolutely reproducible but the quantitative values and particularly the total iron concentration at the end of the experiments, the relative proportion between carbonates and oxides, and the final TIC values exhibit a rather high inaccuracy. Everything seems to indicate that small changes in the experimental conditions lead to significant changes in the subsequent chemical and mineralogical arrangement of the system. This outcome is not strange given the disequilibrium conditions under which the system works and the presence of some critical parameters that can modify the system evolution. On the other hand, the need for a better knowledge of the magnetite-magnesioferrite solid solution, its solubility and mixing properties is obvious. The available data for magnesioferrite solubility are not suitable for temperatures higher than 40 ºC. Anyhow, SIMfr < SIMag < SIHem throughout the experiments and the available data indicate that the precipitation of Mfr was negligible (see next section). While determining the Fe(II)/Fe(III) ratio in the precipitate is feasible using state-of-the-art techniques, monitoring the evolution of this ratio in the aqueous solution is not an analytical problem but a problem related to the kinetics and reproducibility of these types of reactions.

4.6. A troublesome assumption in modeling water-mineral interactions

Irreversible reactions are frequently described by a sequence of states of “partial equilibrium” between the aqueous solution and the successive phases produced along the reaction path.46 In AQ21



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Page 22 of 30

SS systems a typical case of partial equilibrium occurs when an initial solid-solution crystal becomes isolated from the aqueous solution by an outer layer of different composition at equilibrium with the aqueous solution. Surface passivation during the uptake of toxic ions by coprecipitation on mineral surfaces is a glaring example.20 In turn, when growth of solid solutions occurs near equilibrium, the process can be interpreted as a sequence of partial equilibrium states.47 Partial equilibrium implies equilibrium with respect to one or more processes in a given system, but disequilibrium with respect to others.46 For example, non-equilibrium phenomena are recognized to be persistent in mineral growth and mineral-water interaction processes. However, most geochemical calculations assume that when a solid dissolves/grows and releases/uptakes solutes to the aqueous phase, the distribution of aqueous species (free ions and complexes) becomes instantaneously re-equilibrated. Such simplifying hypothesis is useful but troublesome, particularly in the case of species involved in redox reactions.42,44 For example, aqueous Fe(II)-Fe(III) oxidation can occur at a slower rate than precipitation or dissolution of Fe-bearing minerals. Here, the coupling between the kinetics of both processes is a main source of uncertainty. In PHREEQC and other speciation programs, the proportion Fe(II)/Fe(III) and the activities of the corresponding aqueous ions are adjusted to redox equilibrium and the saturation index with respect to Fe(III)bearing phases is determined from these values. While this approach is suitable in predicting equilibrium end-points, does not provide saturation indices in “real time” since some reactions between aqueous species are actually not instantaneous and can be even slower than the dissolutionprecipitation of related solids. The chemical (Fe/Mg) and mineralogical (carbonate/oxide minerals) distribution at the end of the experiments can be estimated by mass balance, which involve (among other calculations) subtraction between initial and final aqueous concentrations of Mg, Fe, and TIC with the subsequent error propagation (the result of such subtractions will be a smaller number with a relatively larger uncertainty). Figure 10 displays the distribution of Mg and Fe in the solid phases (obtained in 228 g of water) as a function of time (after the initial precipitation event) in the case of the experiments D3-D6. Despite the inaccuracy, there is a good qualitative matching with the previous observations. As shown in Figure 10a, the total amount of iron decreases (dissolution) in carbonates and increases (precipitation) in oxides as the reaction time passes by. Moreover, the amount of iron in (Mg,Fe)CO3 with composition XSid-2 is significantly lower than the amount in (Mg,Fe)CO3 with composition XSid-1.

22


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Figure 10. Distribution of Mg and Fe in the solid phases as a function of time (mmoles obtained using 228 g of water). Experiments D3-D6. Reaction temperature = 150 ºC; initial Fe/Mg ratio = 1. (a) Millimoles of iron in carbonates and oxides. (b) Millimoles of magnesium in carbonates. (c) Millimoles of carbonates and oxides as a function of time. Magnesium behaves in the opposite way (Fig. 10b). Firstly, the incorporation of magnesium into magnetite and other oxide phases seem to be negligible since mass balance calculations give indifferently positive and negative values within the error range. Moreover, the amount of Mg in (Mg,Fe)CO3 of composition XSid-2 is initially lower than the amount in (Mg,Fe)CO3 with composition XSid-1, but the total amount of solid solution increases (precipitation) with the reaction time. Finally, as the reaction time passes by (Fig. 10c) all Mg-in-solids accumulates in the carbonate phase while iron concentrates in oxides: The equilibrium end-point would be a solid solution of composition XMgs > 0.99 at equilibrium with both hematite and the aqueous solution. In all the cases the maximum uncertainty corresponds to ~7 days of reaction, just when dissolution of (Mg,Fe)CO3 with composition XSid-1 is most actively coupled with precipitation of Mg-rich (Mg,Fe)CO3. The results displayed in Figure 10 are consistent with the average Fe2+/Fe3+ ratio estimated by XPS. As a matter of example, Figure 11 shows the Fe2p region of the XPS spectrum of the precipitate D5 (taken after 10 days of reaction). As can be observed, about 19.3 mole% iron occurs 23



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Page 24 of 30

as Fe3+, whereas the calculations represented in Figure 10 indicate that about 18.8 mole% iron occurs as oxy-hydroxides. We have obtained similar XPS results in other precipitates of the series C and D. The fact that the amount of iron in oxy-hydroxides is of the same order than the total amount of Fe3+ seems to indicate that the magnetite proportion in the non-carbonate minerals is not very high. However, we should be cautious in reaching conclusions because the error in the solid compositions determined by mass balance is relatively high (>10%).

Figure 11. Fe2p region of the XPS spectrum of the precipitate D5. Table 3. Equilibrium assemblage of phases (mmoles obtained using 228 g of water) Exp.

Fe in oxides (mmoles)

(Mg,Fe)CO3 (mmoles)

XSid 24


Page 25 of 30

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A

2.25

20.2

0.0014

B

4.51

17.9

0.0029

C

9.04

13.3

0.0061

D

11.3

11.0

0.0077

E

11.4

8.63

0.0094

F

18.2

3.98

0.0127

G

20.5

1.62

0.0145

5. CONCLUDING REMARKS AND FUTURE WORK

Determining equilibrium end-points of mineral-water reactions using a suitable speciation code like PHREEQC is straightforward. In the case of the experiments reported in this work, equilibrium involves two solid phases, hematite and a magnesium-rich solid solution, as shown in Fig. 10c and Table 3. The assemblage of solid phases is quite the same in all the experiments compiled in Table 1: Despite the nitrogen atmosphere, most iron will tend to accumulate in hematite after Fe(II)→Fe(III) oxidation of the aqueous ions. However, reaching such an end-point takes a long time. In Fe(II)-bearing carbonated waters, the initial precipitation of (Mg,Fe)CO3 is kinetically fc favored and Fe(II) accumulates in metastable members (richer in Mg than ) of this solid

solution. As a consequence, the initial precipitate undergoes a dissolution-recrystallization process driven by two opposite tendencies: The strong tendency of the aqueous Fe2+ ions to partition towards the solid solution and the oxidation subsequent precipitation of the ferrous ions. Fe(II)→Fe(III) oxidation is a slow process, particularly in our experimental scenario, due to the presence of high concentrations of chlorine and TIC. The subsequent removal of Fe(III) occurs mainly by crystallization of Mag-Mfr with very minor amounts of magnesium and amorphous oxyhydroxides. Magnetite is also metastable since the aqueous solution is maximally supersaturated with respect to hematite, at least if we assume that the equilibrium distribution of aqueous species is attained instantaneously. Anyhow, we can still keep this “safe” assumption44 and “charge the bill” to the favorable nucleation kinetics of magnetite. According to the previous discussion, the reaction path to equilibrium involves dissolution, recrystallization and oxidation processes whose combined kinetics is difficult to predict, especially in the presence of partial equilibrium scenarios (surface passivation, concentric zoning, etc.). From the point of view of CO2 sequestration by injection in basaltic rocks, the good news are that most magnesium is accumulated in carbonate phases. Basaltic olivines are Mg-rich but the fact that iron 25



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Page 26 of 30

tends to accumulate in non-carbonate minerals could decrease the yield of the storage process. The redox conditions during in situ carbonation should be accurately determined to guarantee that iron could take part to the storage process. The next challenge would be introducing calcium in the ‘cocktail’. The classical division of olivine minerals into olivines that contain virtually no calcium, (Fe,Mg)2SiO4 and olivines with calcium Ca(Fe,Mg)SiO4 reflects the limited substitution of Ca2+ (~1Å) for Fe2+ (~0.78Å) or Mg2+ (~0.72Å) in the structure of ferromagnesian silicates (olivines, piroxenes, amphiboles, etc.). In contrast, the ability of Fe2+ to substitute for Mg2+ favors the development of complete solid solution series in ferromagnesian silicates. A question arises: can we classify the products of carbonation in a similar way, i.e. ferromagnesian carbonates with and without calcium? The answer is affirmative. Natural occurrences of carbonate minerals in the ternary system CaCO3-MgCO3-FeCO3 can belong to either the ankerite-dolomite Ca(Fe,Mg)(CO3)2 or the siderite-magnesite (Fe,Mg)CO3 series. Intergrowth between members of both solid solutions are frequent not only in terrestrial materials,48 but also in carbonaceous meteorites.49 Interestingly, the partitioning of Mg2+ and Fe2+ between cocrystallizing Ca(Fe,Mg)(CO3)2 and (Fe,Mg)CO3 depends on temperature and has been proposed as a geothermometer.15,50 Assessing the extent to which the redox and kinetic effects described in this paper affect the crystallization behavior in the CaCO3-MgCO3-FeCO3 system is a main goal for future work in this field.

ACKNOWLEDGEMENTS

This work was supported by Marie Curie Initial Training Network “Geologic Carbon Storage” (European Comission, FP7-People-ITN-CO2-REACT-317235). XRD, ICP-MS, SEM-EDS, and XPS analyses were carried out at the Scientific-Technical Services of the University of Oviedo.

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For Table of Contents Use Only

Dissolution-recrystallization of (Mg,Fe)CO3 during hydrothermal cycles: FeII/FeIII conundrums in the carbonation of ferromagnesian minerals Fulvio Di Lorenzo and Manuel Prieto Department of Geology, C/Jesús Arias de Velasco s/n, University of Oviedo, 33005 Oviedo, Spain

Synopsis During hydrothermal heating, metastable (Mg,Fe)CO3 precipitates obtained at high supersaturation undergo a dissolution-recrystallization process driven by two forces: The preferential partition of Fe2+ towards (Mg,Fe)CO3 and the oxidation of aqueous Fe2+ species and subsequent precipitation of Fe(III)-bearing solids. Considering the coupling between both forces is essential in modeling the artificial carbonation of ferromagnesian minerals for CO2 sequestration purposes.

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