Do organic ligands affect forsterite dissolution rates?

Applied Geochemistry 39 (2013) 69–77

Contents lists available at ScienceDirect

Applied Geochemistry journal homepage: www.elsevier.com/locate/apgeochem

Do organic ligands affect forsterite dissolution rates? Julien Declercq, Olivier Bosc, Eric H. Oelkers ⇑ GET/Université Paul Sabatier, Observatoire Midi-Pyrénées, UMR 5563 (CNRS/UPS/IRD/CNES), 14 Avenue Edouard Belin, 31400 Toulouse, France

a r t i c l e

i n f o

a b s t r a c t

Article history: Received 8 November 2012 Accepted 27 September 2013 Available online 4 October 2013 Editorial handling by M. Hodson

Far-from equilibrium, steady state forsterite dissolution rates were measured at pH 3 and 25 °C in aqueous solutions containing 0.1 m/kg NaCl and up to 0.1 mol/kg of 13 distinct dissolved organic ligands in mixed-flow reactors. The organic ligands considered in this study include those common in Earth surface environments and those considered as potential catalysts for use in CO2 sequestration efforts: acetate, oxalate, citrate, EDTA4, glutamate, gluconate, malonate, aspartate, tartrate, malate, alginate, salycilate and humate. The presence of up to 0.1 mol/kg of each organic ligand altered forsterite dissolution rates less than 0.2 log units, which is the estimated uncertainty of the measured rates. Results obtained in this study, therefore, suggest that the presence of aqueous organic anions negligibly affects forsterite far-from equilibrium dissolution rates in most natural environments, and indicate that forsterite carbonation may not be appreciably accelerated by organic ligand catalysis. Ó 2013 Elsevier Ltd. All rights reserved.

1. Introduction A significant number of studies have been aimed at characterizing forsterite dissolution rates at various solution compositions and temperatures (Luce et al., 1972; Sanemasa et al., 1972; Grandstaff, 1978, 1986; Murphy and Helgeson, 1987, 1989; Blum and Lasaga, 1988; Van Herk et al., 1989; Wogelius and Walther, 1991, 1992; Casey and Westrich, 1992; Jonckbloedt, 1998; Awad et al., 2000; Chen and Brantley, 2000; Rosso and Rimstidt, 2000; Pokrovsky and Schott, 1999, 2000a,b; Oelkers, 2001; Giammar et al., 2005; Golubev et al., 2006; Hänchen et al., 2006; Olsen and Rimstidt, 2008; Rimstidt et al., 2012; Saldi et al., 2013; Wang and Gaimmar, 2013). There has been increased recent interest in these rates owing to the potential application of forsterite in carbon capture and storage efforts (e.g. Giammar et al., 2005; Oelkers and Schott, 2005; Marini, 2007; Kelemen and Matter, 2008; Oelkers et al., 2008; Dufaud et al., 2009; Prigiobbe et al., 2009; Garcia et al., 2010; King et al., 2010a,b; Daval et al., 2011; Guyot et al., 2011; Olsson et al., 2012). Forsterite is commonly thought of as the best source of the divalent metal cations required for mineral carbonation due to its fast dissolution rates and global abundance. The carbonation of forsterite in CO2-rich fluids can occur via

Mg2 SiO4ðsÞ þ2CO2 þ 2H2 O ¼ 2MgCO3ðsÞ þH4 SiO4ðaqÞ Forsterite

ð1Þ

Magnesite

where the resulting aqueous silica could eventually precipitate as amorphous silica or other silicate minerals (cf. Weres et al., 1981; Teir et al., 2007; Orlando et al., 2001). Numerous studies have focused on the application of reaction (1) to carbon sequestration ⇑ Corresponding author. Tel.: +33 561332575. E-mail address: [email protected] (E.H. Oelkers). 0883-2927/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.apgeochem.2013.09.020

on an industrial scale (e.g. O’Connor et al., 2000a,b; Wolf et al., 2004; Giammar et al., 2005; Maroto-Valer et al., 2005; Chen et al., 2006; Gerdemann et al., 2007; Oelkers et al., 2008; Broecker, 2012). Reaction (1) involves the coupling of two processes: forsterite dissolution and magnesite precipitation. The overall rate of carbonation reaction (1) is, therefore, a function of these two processes (e.g. Saldi et al., 2013). Catalysis which could accelerate the rates of forsterite dissolution thus may also accelerate its carbonation in accord with reaction (1). A vast number of past studies have suggested that organic ligands can increase the dissolution rates of silicate minerals substantially (e.g. Huang and Kiang, 1972; Manley and Evans, 1986; Mast and Drever, 1987; Bennett et al., 1988; Welch and Ullman, 1993, 1996; Poulson et al., 1997; Cama and Ganor, 2006; Golubev and Pokrovsky, 2006; Golubev et al., 2006; Ganor et al., 2009; Pokrovsky et al., 2009; Schott et al., 2009). The effect of the presence of organic ligands on forsterite dissolution has been considered in detail by Grandstaff (1986), Wogelius and Walther (1991), Hänchen et al. (2006), Olsen and Rimstidt (2008) and Prigiobbe and Mazzotti (2011). Grandstaff (1986) reported an increase of forsterite dissolution rates when the reactive aqueous fluid contained citrate, EDTA, oxalate, tannic acid, succinate and phthalate at 25 °C and pH 3.5 and 4.5. Wogelius and Walther (1991) reported that the addition of either 0.05 mol/kg potassium phthalate or 103 mol/kg ascorbic acid increases forsterite dissolution rates at 25 °C by 0.75 log units at pH 4, but this effect decreased with decreasing pH. Hänchen et al. (2006) reported that 103 mol/kg citric acid increased forsterite dissolution rates at 90 °C by 0.25 log units at pH 3.4 and 0.5 log units at pH 4.5. Olsen and Rimstidt (2008) reported that 104 mol/kg oxalic acid increases forsterite dissolution rates by a factor of 6 at 25 °C and 2.5 < pH < 6.5.

70

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

There have been a number of distinct interpretations of the observed effect of organic ligands on silicate dissolution rates. Stumm and collaborators (Furrer and Stumm, 1986; Stumm and Wollast, 1990) proposed that a parallel organic ligand promoted dissolution pathway catalyzes the formation of organic anion surface species leading to enhanced rates. Oelkers and Schott (1998) proposed that organic ligands enhance silicate dissolution rates by forming aqueous complexes with those aqueous metal cations that would otherwise slow rates. Yet others have suggested the formation of surface chelates by organic ligands as a rate promoting mechanism (e.g. Wogelius and Walther, 1991). This study aims to further understanding of organic ligands on mineral reactivity by measuring the effect of the presence of organic ligands on forsterite dissolution kinetics at 25 °C and pH 3. The ligands investigated included acetate, oxalate, citrate, EDTA4, glutamate, gluconate, malonate, aspartate, tartrate, malate, alginate, salycilate and humate. The purpose of this paper is to report the results of this experimental study and to use these results to further illuminate the effect of the presence of organic ligands on forsterite dissolution rates in natural and carbon storage processes.

2. Materials and methods 2.1. Materials Xenoliths containing forsteritic (Fo89) olivine were collected from the San Carlos volcanic field east of Globe, AZ, USA. The xenoliths were crushed and forsterite was recovered by handpicking. This forsterite was then ground and sieved to obtain the 100– 200 lm size fraction. X-ray diffraction analysis of this material revealed it to be pure forsterite, free of clays and secondary phases. Following the protocol established by Olsen and Rimstidt (2008), the ground forsterite was first ultrasonically cleaned in 0.01 M HCl for 10 min, left to settle for 5 min, and washed with alcohol to remove the fine particles. The resulting forsterite was then dried and used in the experiments without further treatment. The BET surface area was measured by N2 adsorption using an Absorb-1 Quantachrome™ and calculated to be 980 cm2/g (±10%). Microscopic analysis of fresh and altered forsterite surfaces was performed using a Jeol JSM840a scanning electron microscope. A photomicrograph of the forsterite prior to the experiments is shown in Fig. 1a. The surfaces are free of both fine particles and secondary phases. Forsterite was dissolved in mixed flow reactors in fluids prepared from 18 MX ultrapure water (MilliQ Plus system), 0.1 mol/

Fig. 1. SEM images of the forsterite powder. Image A shows the forsterite before dissolution experiments while image B shows the forsterite powder after it had been dissolved during experimental series 6. Note the initial forsterite is free of fine particles, and appears to have sharp edges while the reacted mineral powder exhibits rounded edges.

kg NaCl, sufficient HCl or NaOH to adjust the pH to 3, and from 0.0001 to 0.1 M of acidic acid, apartic acid, citric acid, EDTA disodium salt dehydrate, sodium gluconate, L-glutamic acid, malic acid, malonic acid, oxalic acid, salicylic acid, tartaric acid, or humic acid. All chemicals used in the preparation of inlet fluids were analytical grade and purchased from either Fluka, Sigma, or Sigma–Aldrich. The organic acids or salts were used as purchased except for the humic acid, which was pre-dissolved using techniques reported by Pokrovsky et al. (2005). A list of the acids or salts used in this study and their chemical formulas is provided in Table 1.

Table 1 List of the dissociation constants and chemical formula for the organic ligands and the identity of the acid or salt used to make the inlet fluids used in this study. With the exception of humate, dissociation constants were compiled from Perrin et al. (1981) and Ullman and Welch (2002). Ligand

Formula

pKaa

Salt or acid used

Salt formula

Acetate Alginate Aspartate

C2 H3 O 2 C2 H3 ðNH2 ÞðCOOÞ2 2

4.76 3.5 2.0–3.9

Acidic acid Alginic acid Aspartic acid

CH3COOH (C6H8O6)n C2H3(NH2)(COOH)2

Citrate

C6 H5 O3 7

3.14–4.75–6.40

Citric acid

C6H8O7

C10 H12 N2 O4 8 C6 H11 O 7

2.0–2.7–6.2–10.3

EDTA disodium salt dihydrate

C10H14N2Na2O8

Sodium gluconate L-Glutamic

C6H11NaO7 C3H5(NH2)(COOH)2

EDTA4

(C6H8O6)n

Gluconate Glutamate

C3 H5 ðNH2 ÞðCOOÞ2 2

3.86 2.1–4.1

Malate

C4 H4 O2 5

3.4–5.2

Malic acid

C4H6O5

Malonate

CH2 ðCOOÞ2 2

2.85–5.70

Malonic acid

CH2(COOH)2

Oxalate

C2 O2 4 C7 H5 O 3

1.2–4.2

Oxalic acid

C2H2O4

2.97 3.0–4.3

Salicylic acid Tartaric acid

C7H6O3 C4H6O6

Salicylate Tartrate

C4 H4 O2 6

acid

Data for humic acid is not available. a Where more than one value is given the values refer to the equilibrium constants sequentially for the first, second, third, and fourth deprotonation reaction where relevant.

71

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77 Table 2 Results of all experiments performed in the absence of aqueous organic ligands. Exp

cSi (mol/kg)  105

cMg (mol/kg)  105

cMg/cSi

pH

Flow (mL/min)

m (g)

log r+,Si (mol/cm2/s)

log r+,Mg (mol/cm2/s)

1.s 1.f 2.s 2.f 3.s 3.f 4.s 4.f 5.s 5.f 6.s 6.f 7.s 7.f 8.s 8.f 9.s 9.f 10.s 10.f 11.s 11.f 12.s 12.f 13.s 13.f

12.9 15.1 11.1 7.57 17.2 4.46 14.5 3.60 20.7 4.75 3.28 1.87 14.7 4.82 14.8 4.68 9.27 1.90 4.29 7.72 4.66 5.02 1.93 1.61 1.64 2.25

26.21 31.01 20.51 15.35 31.44 8.35 27.01 7.21 33.72 8.61 6.75 4.29 26.02 9.42 26.09 8.06 10.04 3.56 8.67 13.39 9.09 8.86 4.14 2.86 3.21 4.10

2.03 2.05 1.85 2.03 1.83 1.87 1.86 2.00 1.63 1.81 2.06 2.29 1.77 1.95 1.76 1.72 1.08 1.87 2.02 1.73 1.95 1.76 2.15 1.78 1.95 1.82

3.33 3.37 3.42 3.05 3.34 2.97 3.04 3.00 3.38 3.01 3.12 3.11 3.06 3.08 3.49 3.04 3.24 3.09 3.04 3.00 3.03 2.90 2.90 2.90 2.90 2.97

0.14 0.14 0.10 0.10 0.15 0.15 0.12 0.14 0.15 0.16 0.40 0.40 0.22 0.22 0.14 0.10 0.15 0.14 0.15 0.15 0.20 0.20 0.30 0.30 0.17 0.17

0.49 0.49 0.31 0.31 0.31 0.31 0.30 0.30 0.30 0.30 0.31 0.31 0.50 0.50 0.31 0.31 0.33 0.33 0.31 0.31 0.30 0.30 0.35 0.35 0.35 0.35

8.65 8.58 8.66 8.82 8.29 8.87 8.45 8.98 8.20 8.80 8.58 8.83 8.41 8.89 8.59 9.03 8.56 9.32 8.88 8.62 8.72 8.69 8.93 9.01 9.25 9.11

8.60 8.53 8.65 8.78 8.28 8.86 8.44 8.94 8.24 8.80 8.53 8.73 8.40 8.93 8.60 9.05 8.78 9.31 8.83 8.64 8.69 8.70 8.86 8.96 9.21 9.11

X.s and X.f designates the first and last experiments of the series.

2.2. Experimental methods All experiments were performed in 30 mL polypropylene mixed-flow reactors (cf. Chaïrat et al., 2007; Flaathen et al., 2010); schematic illustrations of the reactor design have been presented by Kohler et al. (2005). The reactors were continuously stirred with floating Teflon stirring bars. These reactors were immersed in a water bath held at a 25 ± 2 °C. The fluid was injected using a Gilson peristaltic pump, which allows fluid flow rates from 0.01 to 10 g/min. The fluid left the reactor through a 0.45 lm acetate filter. No additional filtering was performed on outlet fluid samples prior to chemical analysis. Inlet and outlet fluid pH was measured using a combination glass electrode calibrated with NIST buffers (pH = 4.002, 6.865, and 9.180 at 25 °C). The precision of pH measurements was ±0.02 units. Magnesium concentrations were determined by flame atomic absorption spectrophotometry using a Perkin Elmer 5100 PC spectrometer equipped with an AS-90 autosampler, with an uncertainty of ±2%. Silica concentrations were determined using the molybdate blue colorimetric method (Govett, 1961) using a Technicon analyzer with an uncertainty of ±2% for most fluids, though some interference was evident in some organic ligand-rich fluids, see below. Seventy-seven mixed-flow reactor experiments were run at a pH of 3 ± 0.4 throughout the study. To allow direct comparison between rates in the presence and absence of organic ligands, experiments were organized into thirteen series. Each series is designated by the prefix of the experiment. For example, series 2 consists of experiments 2-b, 2-1, 2-2, 2-3, 2-4 and 2-f. Each experimental series was initiated by passing an organic-free, pH 3 fluid through the reactor at a constant flow rate for approximately 6 days. Once an initial steady state was obtained and confirmed, the original inlet fluid was replaced by a pH 3 fluid containing 104 mol/kg of a selected organic ligand until a second steady state was obtained and confirmed. This sequence was generally repeated using pH 3 fluids containing 103, 102, 101 and 0 mol/kg of this organic ligand. The total run time for each complete experimental series ranged from 3 to 6 weeks. By running the experiments in

series dissolution rates measured in the presence of organic ligands can be compared directly to rates measured in organic ligand-free fluids on the identical forsterite samples. After the steady-state was obtained for the final organic ligand-free fluid, the reactor was dismantled, the remaining forsterite powder was recovered and the reactor cleaned. In the case of EDTA and L-glutamate, organic ligand concentrations of 0.1 M could not be achieved due to solubility constraints. Approximately 10 mL of outlet fluid was collected twice daily during each experiment. Steady-state was confirmed by the outlet fluids maintaining a constant Mg and Si concentration within analytical uncertainty for a minimum of 3 residence times. The residence time is defined as the volume of the reactor divided by the reactive fluid flow rate, which was fixed at 0.15–0.40 g/min in the experiments. Further details of each experimental series are provided in Tables 2 and 3. 3. Results An example of the reactive fluid Mg and Si concentrations measured during one experimental series is depicted in Fig. 2. This figure shows the reactive fluid evolution during experiments designed to assess the effect of aqueous malonate on forsterite dissolution rates. It can be seen in this figure that the addition of the organic ligand leads to the attainment of a new steady state condition as indicated by the dashed lines in the figure. Forsterite steady-state dissolution rates were calculated from the aqueous Mg and Si concentrations using the relationship:

rþ;i ¼

qðc i;out  c i;in Þ S  m  mi

ð2Þ

where q refers to the fluid flow rate, ci,in and ci,out represent the aqueous concentration of the ith element in the inlet and outlet, S denotes the specific surface area, m designates the amount of forsterite powder introduced in the reactor at the beginning of the experiment series and mi stands for the stoichiometric number of moles of the ith element in one mole of dissolving forsterite.

72

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

Table 3 Results of all experiments performed in the presence of aqueous organic ligands. EXP

Ligand

Ligand conc. (mol/kg)

cSi (mol/kg)  105

cMg (mol/kg)  105

cMg/cSi

pH

Flow (mL/min)

m (g)

log r+,Si (mol/cm2/s)

log r+,Mg (mol/cm2/s)

1-1 1-2 1-3 1-4 1-5

Malonate

0.0001 0.0001 0.001 0.01 0.1

15.8 6.50 5.18 5.39 6.14

28.9 12.2 9.13 9.13 12.2

1.83 1.87 1.76 1.67 1.98

3.38 3.39 3.14 3.01 3.01

0.14 0.16 0.16 0.16 0.16

0.49 0.30 0.30 0.30 0.30

8.55 8.67 8.77 8.76 8.70

8.54 8.66 8.78 8.78 8.66

2-1 2-2 2-3 2-4

Acetate

0.0001 0.001 0.01 0.1

14.9 15.3 13.5 8.36

26.9 28.5 25.1 24.5

1.81 1.86 1.85 2.93

3.50 3.22 3.17 3.05

0.10 0.10 0.10 0.10

0.31 0.31 0.31 0.31

8.51 8.50 8.55 8.76

8.51 8.49 8.54 8.55

3-1 3-2 3-4 3-5

Oxalate

0.0001 0.001 0.01 0.1

11.2 8.37 8.45 1.42

21.8 16.4 16.7 5.80

1.94 1.96 1.98 4.09

3.34 3.11 2.95 3.03

0.30 0.30 0.30 0.30

0.31 0.31 0.31 0.31

8.17 8.30 8.29 9.07

8.14 8.26 8.26 8.71

4-1 4-2 4-3 4-4

Citrate

0.0001 0.001 0.01 0.1

6.86 5.75 6.76 4.12

12.1 10.9 13.2 n.d.

1.77 1.90 1.96 n.d.

3.17 3.13 3.04 3.07

0.30 0.30 0.30 0.30

0.30 0.30 0.30 0.30

8.38 8.46 8.39 8.60

8.39 8.44 8.35 n.d.

5-1 5-2 5-3

EDTA

0.0001 0.001 0.01

11.0 5.71 3.82

22.9 11.2 9.01

2.09 1.96 2.36

2.95 3.10 3.20

0.30 0.30 0.30

0.30 0.30 0.30

8.17 8.45 8.63

8.11 8.42 8.51

6-1 6-3 7-1 7-2 7-3 7-4 7-5

Tartrate

0.001 0.01 0.0001 0.0001 0.001 0.01 0.1

3.26 3.00 4.88 5.11 4.63 4.93 12.4

6.72 5.94 9.23 9.79 8.52 8.85 24.1

2.06 1.98 1.89 1.92 1.84 1.79 1.94

3.07 3.20 3.12 3.13 2.99 2.99 3.08

0.39 0.40 0.22 0.30 0.30 0.30 0.14

0.31 0.31 0.30 0.30 0.30 0.30 0.50

8.60 8.62 8.67 8.51 8.55 8.53 8.67

8.54 8.58 8.65 8.65 8.55 8.53 8.65

Malate

8-1 8-2 8-3 8-4

Salicylate

0.0001 0.001 0.01 0.1

15.0 4.11 4.93 5.18

27.4 7.45 8.43 44.3

1.83 1.81 1.71 8.72

3.01 3.03 3.01 3.11

0.14 0.10 0.10 0.10

0.50 0.31 0.31 0.31

8.59 9.08 9.01 8.98

8.58 9.08 9.03 8.30

9-1 9-2 9-3 9-4

Alginate

0.0001 0.001 0.01 0.1

4.69 1.27 2.08 1.64

9.53 2.61 3.84 3.41

2.03 2.05 1.85 2.08

3.15 3.29 3.30 3.04

0.15 0.31 0.31 0.31

0.30 0.31 0.31 0.31

8.85 9.11 8.90 9.00

8.80 9.05 8.89 8.94

10-1 10-2 10-3 10-4 10-5

Humatea

0.0002 0.0006 0.0006 0.001 0.001

3.43 2.92 3.24 3.70 14.5

6.28 5.43 5.08 6.83 n.d.

1.83 1.86 1.57 1.60 n.d.

3.07 3.04 3.05 3.04 3.22

0.33 0.30 0.15 0.30 0.10

0.14 0.31 0.33 0.31 0.33

8.30 8.76 9.05 8.65 8.57

8.29 8.75 9.11 8.65 n.d.

11-1 11-2 11-3 11-4

Gluconate

0.0001 0.001 0.01 0.1

4.36 1.91 2.00 3.46

9.83 4.52 4.10 6.79

2.26 2.36 2.05 1.96

3.15 3.51 3.43 3.15

0.15 0.20 0.20 0.20

0.30 0.30 0.30 0.30

9.70 9.11 9.09 8.85

8.79 8.99 9.04 8.82

13-1 13-2 13-3

Glutamate

0.0001 0.001 0.01

1.65 1.86 1.93

3.29 3.81 3.76

1.99 2.05 1.95

2.77 2.88 2.81

0.21 0.21 0.21

0.35 0.35 0.35

9.21 9.16 9.15

9.17 9.11 9.11

12-1 12-2 12-3 12-4

Aspartate

0.0001 0.001 0.01 0.1

2.20 2.08 2.00 1.80

5.07 5.07 5.07 5.07

2.30 2.43 2.53 2.82

2.92 2.95 2.92 2.92

0.20 0.20 0.20 0.20

0.35 0.35 0.35 0.35

9.11 9.14 9.15 9.20

9.01 9.01 9.01 9.01

n.d. Not determined. a The ligand concentration for the humates is given in units of volume fraction.

Steady-state Si and Mg outlet concentrations and the dissolution rates calculated from all experiments are provided in Tables 2 and 3. With few exceptions the molar Mg to Si ratio of all steady-state outlet flow is 1.78 ± 0.25, consistent with stoichiometric dissolution of the forsterite. The logarithm of the dissolution rates calculated based on Si release is plotted as a function of those based on Mg release in Fig. 3. Only 3 experiments are significantly non-stoichiometric: 11-1, 3-4 and 8-4, containing respectively 104 mol/kg of gluconate, and 0.1 mol/kg of oxalate and salicylate. Tests of the molybdate blue method used to analyze dissolved Si in this study show that the presence of 0.1 mol/kg gluconate and oxalate significantly affected Si concentrations measured by this

method; these observations will be expanded upon in detail in a future communication. As such, rates based on Mg release are used for the interpretation of these experiments in the discussion below. All experimentally determined forsterite dissolution rates are plotted as a function of organic ligand concentration in Figs. 4 and 5. Plots in these figures are grouped by experimental series allowing the direct assessment of the effect on forsterite dissolution rates of each organic ligand. Measured rates are plotted in each figure in the sequence they were measured; first rates were measured in an organic ligand-free fluid, then rates were measured in a sequence of fluids containing increasing organic ligand

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

73

Fig. 2. The temporal evolution of reactive fluid aqueous Si and Mg concentrations during experimental series 1. The solid symbols correspond to measured concentrations, the vertical solid lines indicate times when the inlet fluid composition was changed, and the dashed lines indicate the steady state concentration values for each experiment.

release, is as much as 0.6 log units different from that measured in the first organic ligand-free fluid. This difference is not systematic; in some experiments the final steady state forsterite dissolution rate is higher and in some the final rate is lower than the initial rate. The origin of these variations remains unclear, but a number of previous studies have suggested that significant temporal variations in rates during experiments are due to changes in reactive surface area (Lüttge et al., 1999; Gautier et al., 2001; Hodson, 2006; Lüttge, 2006); mineral dissolution rates are commonly assumed to be proportional to the mineral–fluid reactive surface area, yet the variation of these surface areas as minerals dissolve is still poorly defined (e.g. Noiriel et al., 2009). A second relevant observation is that rates measured in the presence of organic ligands are equal to those measured in organic ligand-free fluids within analytical uncertainty.

4. Discussion 4.1. Comparison with past studies Fig. 3. Measured far-from-equilibrium steady-state forsterite dissolution rates at 25 °C and pH 3 based on Si release as a function of corresponding rates based on Mg release. Rates measured in the presence of aqueous organic ligands are depicted as solid squares were whereas those measured in organic ligand-free fluids are depicted as solid diamonds. Unfilled symbols correspond to experiments showing an apparent non-stochiometric metal release. Note that 74 of the 77 plotted rates are within 0.25 log unit of the 1:1 stoichiometric line.

concentrations and, finally, a second rate measurement was performed in an organic ligand-free fluid. First it is evident that the final rate measured in organic ligand-free fluids, based on Mg

Wogelius and Walther (1991) reported that the degree to which ascorbic acid and K-phthalate affect forsterite dissolution rates depends on pH. They observed no effect of these organic compounds at pH 2, but reported that these ligands increased forsterite dissolution rates by up to 0.75 log units at pH 6. Grandstaff (1986) noted a similar pH dependence of the effect of organic ligands on forsterite dissolution rates. Hänchen et al. (2006) reported that the effect of citric acid on forsterite dissolution rate was small under acid conditions but increases with increasing pH at 90 °C. In contrast, Olsen and Rimstidt (2008) suggested that the presence of oxalate on forsterite dissolution rates at 25 °C was pH independent from 2.5 < pH < 6.5. One

74

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

Fig. 4. Forsterite far from equilibrium dissolution rate conditions at 25 °C and pH 3 measured in the present study as a function of the concentration of the indicated organic ligand concentration. Rates obtained in the presence of aqueous organic ligands are shown as solid diamonds. Rates obtained at the beginning and end of the corresponding experimental series in organic ligand-free fluids are plotted as filled triangles at a log organic ligand concentration of 5 and 0, respectively, to illuminate the effect of the ligands on rates.

possible reason why this latter study is inconsistent with the previous studies is that their conclusions were based on forsterite dissolution experiments that lasted only 2 h. As can be seen in Fig. 2, the Mg and Si release rates during the first few hours

of forsterite dissolution experiments can be inconsistent with the longer-term steady state rates. Rates in the present study demonstrate that at pH 3 forsterite dissolution rates are negligibly impacted by the presence of a wide

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

75

Hood, 1965; Suess, 1970; Otsuki and Wetzel, 1972; Reddy, 1977; Berner et al., 1978; Reynolds, 1978; Reddy and Wang, 1980; Innskeep and Bloom, 1986; Giannimaras and Koutsoukos, 1988; Dove and Hochella, 1993; Katz et al., 1993; Paquette et al., 1996). 4.3. Implications for silicate dissolution models

Fig. 5. Forsterite far from equilibrium dissolution rates at 25 °C and pH 3 measured in the present study as a function of aqueous humate concentration. Rates obtained in the presence of aqueous humate are shown as solid diamonds. Rates obtained at the beginning and end of the experimental series in organic ligand-free fluids are plotted as filled triangles at an aqueous humate concentration of 0 and 0.0011, respectively, to illuminate the effect of the aqueous humate on rates.

variety of organic ligands at 60.1 mol/kg. These observations are in general agreement with the previous reports that aqueous organic ligands have at most a small effect on forsterite dissolution rates in strongly acidic conditions but may have an effect at higher pH. A reason for the contrasting effects of aqueous organic ligands on steady-state forsterite dissolution rates with increasing pH may stem from their aqueous speciation. As can be seen in Table 1, many of the organic ligands considered in this study have a first dissociation constant near pH 3. As such these organic species tend to be present as neutral aqueous species at acidic pH, but as negatively charged aqueous species in mildly acidic and neutral conditions. Another possible reason for the contrasting effect of organic ligands on forsterite dissolution rates with increasing pH is the composition of the forsterite surface. Both the charge and Mg/Si ratio of forsterite surfaces are strong functions of pH, factors which could affect its reactivity (Pokrovsky and Schott, 2000a,b; Oelkers et al., 2009). 4.2. Implications for carbon capture and storage One of the major goals of this study is to assess the degree to which the addition of organic ligands might accelerate Mg release from forsterite thus aiding mineral carbonation efforts. The pH of CO2-charged water depends on the partial pressure of CO2 in the system; the pH of CO2 saturated water decreases from 3.9 to 3.0 with increasing CO2 partial pressure from 1 to 64 bars (Gíslason et al., 2010). The results summarized above suggest that the addition of most organic ligands will have little effect on Mg release from forsterite in these CO2-charged waters. This conclusion agrees with those of Wolff-Boenisch et al. (2006, 2011) who concluded that the addition of various organic ligands negligibly affected the rates of peridotite dissolution at 25 °C and pH 3.6. As forsterite dissolves in CO2 charged waters, however, the fluid will neutralize (e.g. Siever and Woodfort, 1979). Under such conditions, previous studies suggest that some acceleration of Mg release could be attained from the presence of a number of organic ligands. The presence of these organic ligands, however might inhibit the ultimate precipitation of carbonates either through increasing the solubility of these minerals (Innskeep and Bloom, 1986; Lebrón and Suárez, 1998) or by inhibiting their precipitation rates (e.g. Kitano and

There are at least two distinct models for the enhancement of silicate mineral dissolution rates in the presence of organic ligands. A number of studies have argued that organic ligands could adsorb to the mineral surface, providing a new parallel reaction mechanism for the detachment of material from the mineral surface (e.g. Furrer and Stumm, 1986; Amrhein and Suarez, 1988). In contrast, Oelkers and Schott (1998) proposed that the observed acceleration of the dissolution rates of some silicate minerals with organic ligand concentration stems from their ability to make aqueous complexes with aqueous metals that would otherwise inhibit rates. Although the degree to which organic ligands adsorb onto forsterite surfaces has not been measured directly in this study, it seems likely that significant adsorption would occur because at pH 3 forsterite surfaces are positively charged (Pokrovsky and Schott, 2000a; Oelkers et al., 2009) whereas most of the organic species considered in this study are at least partly present as aqueous anions at this pH due to their low dissociation constants (see Table 3). The lack of forsterite dissolution rate acceleration in the presence of most organic ligands seems to be inconsistent with the hypothesis that rates are accelerated by ligand adsorption. Although aqueous Mg-organic ligand complexation does occur, forsterite dissolution rates have been shown to be independent of aqueous Mg2+ activity (Oelkers, 2001). As such this aqueous Mg-organic ligand complexing would not lead to accelerated rates according to the Oelkers and Schott (1998) hypothesis, consistent with the observations reported above. 5. Conclusions The results of this study suggest that the presence of organic ligands at concentrations of <0.1 mol/kg have at most a small effect on forsterite dissolution rates at pH 3 and 25 °C. As this pH is characteristic of the CO2 charged waters that might be found in the subsurface during carbon storage efforts it seems unlikely that the presence of organic ligands would enhance the availability of those divalent metals required for mineral carbonation from ultramafic rocks prior to the neutralization of this fluid. In contrast, the presence of organic ligands have been shown to enhance the dissolution rates of the intermediate feldspars and basaltic glass at acidic conditions (e.g. Blum and Stillings, 1995; Thorseth et al., 1995; Ullman et al., 1996; Oelkers and Schott, 1998; Alt and Mata, 2000; Oelkers and Gíslason, 2001); as such the use of organic ligands might prove effective in accelerating mineral carbonation in basaltic and/or rocks of intermediate composition. Nevertheless, organic ligands may hinder mineral carbonation due to their effect on the solubility and/or precipitation kinetics of carbonate minerals. As such the potential of aqueous organic ligands to enhance carbon mineralization remains questionable. Acknowledgements We would like to thank Alain Castillo for technical assistance throughout the duration of the experimental work, Carole Causserand for her generous help during the analytical part of the work, and Philippe de Parseval for aid creating SEM images. We thank Per Aagaard, Stacey Callahan, Oleg Pokrovsky, Jacques Schott, Pascale Bénézeth and Morgan T. Jones for helpful discussions during the course of this study. Support from Centre National de la

76

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77

Recherche Scientifique, and the European Community through CARB-FIX (Collaborative Project-FP7-ENERGY-2011-1-283148) is gratefully acknowledged.

References Alt, J.C., Mata, P., 2000. On the role of microbes in the alteration of submarine basaltic glass: a TEM study. Earth Planet. Sci. Lett. 181, 301–313. Amrhein, C., Suarez, D.L., 1988. The use of a surface complexation model to describe the kinetics of ligand-promoted dissolution of anorthite. Geochim. Cosmichim. Acta 52, 2785–2793. Awad, A., van Groos, A.F.K., Guggenheim, S., 2000. Forsteritic olivine: effect of crystallographic direction on dissolution kinetics. Geochim. Cosmochim. Acta 64, 1765–1772. Bennett, P.C., Melcer, M.E., Siegel, D.I., Hasset, J.P., 1988. The dissolution of quartz in dilute aqueous solutions of organic acids at 25 °C. Geochim. Cosmochim. Acta 52, 1521–1530. Berner, R.A., Westrich, J.T., Smith, J., Martens, C.S., 1978. Inhibition of aragonite precipitation from supersaturated seawater: a laboratory and field study. Am. J. Sci. 278, 816–837. Blum, A.E., Lasaga, A.C., 1988. Role of surface speciation in the low-temperature dissolution of minerals. Nature 331, 431–433. Blum, A.E., Stillings, L.L., 1995. Chemical weathering of feldspars. Rev. Mineral 31, 173–233. Broecker, W.S., 2012. The carbon cycle and climate change: memoirs of my 60 years in science. Geochem. Perspect. 1, 221–351. Cama, J., Ganor, J., 2006. The effect of organic acids on the dissolution of silicate minerals: a case study of oxalate catalysis of kaolinite dissolution. Geochim. Cosmochim. Acta 70, 2191–2209. Casey, W.H., Westrich, H., 1992. Control of dissolution rates of orthosilicate minerals by divalent metal–oxygen bonds. Nature 355, 157–159. Chaïrat, C., Schott, J., Oelkers, E.H., Lartigue, J.-E., Harouiya, N., 2007. Kinetics and mechanism of natural fluorapatite dissolution at 25 °C and pH from 3 to 12. Geochim. Cosmochim. Acta 71, 5901–5912. Chen, Y., Brantley, S.L., 2000. Forsterite dissolution at 65 °C and 2 < pH < 5. Chem. Geol. 165, 267–281. Chen, Y., Lundqvist, P., Johansson, A., Platell, P., 2006. A comparative study of the carbon dioxide transcritical power cycle compared with an organic rankine cycle with R123 as working fluid in waste heat recovery. Appl. Therm. Eng. 26, 2142–2147. Daval, D., Sissmann, O., Menguy, N., Saldi, G.D., Guyot, F., Martinez, I., Corvisier, J., Garcia, B., Machouk, I., Knauss, K.G., Hellmann, R., 2011. Influence of amorphous silica layer formation on the dissolution rate of olivine at 90 °C and elevated pCO2. Chem. Geol. 284, 193–209. Dove, P.M., Hochella Jr., M.F., 1993. Calcite precipitation mechanisms and inhibition by orthophosphate. In situ observations by scanning force microscopy. Geochim. Cosmochim. Acta 57, 705–714. Dufaud, F., Martinez, I., Shilobreeva, S., 2009. Experimental study of Mg-rich silicates carbonation at 400 and 500 °C and 1 kbar. Chem. Geol. 265, 79–87. Flaathen, T.K., Gislason, S.R., Oelkers, E.H., 2010. The effect of aqueous sulphate on basaltic glass dissolution rates. Chem. Geol. 277, 345–354. Furrer, G., Stumm, W., 1986. The coordination chemistry of weathering: I. Dissolution of d-Al2O3 and BeO. Geochim. Cosmochim. Acta 50, 1847–1860. Ganor, J., Reznik, I.J., Rosenberg, Y.O., 2009. Organics in water–rock interactions. Rev. Mineral. Geochem. 70, 259–369. Garcia, B., Baumont, V., Perfetti, E., Rouchon, V., Blanchet, D., Oger, P., Dromart, G., Huc, A.-Y., Haeseler, F., 2010. Experiments and geochemical modeling of CO2 sequestration by olivine: potential, quantification. Appl. Geochem. 25, 1383– 1396. Gautier, J.-M., Oelkers, E.H., Schott, J., 2001. Are quartz dissolution rates proportional to BET surface areas? Geochim. Cosmochim. Acta 65, 1059–1070. Gerdemann, S.J., O’Connor, W.K., Dahlin, D.C., Penner, L.R., Rush, H., 2007. Ex situ aqueous mineral carbonation. Environ. Sci. Technol. 41, 2587–2593. Giammar, D.E., Bruant, R.G., Peters, C.A., 2005. Forsterite dissolution and magnesite precipitation at conditions relevant for deep saline aquifer storage and sequestration of carbon dioxide. Chem. Geol. 217, 257–276. Giannimaras, E.K., Koutsoukos, P.G., 1988. Precipitation of calcium carbonate in aqueous solutions in the presence of oxalate anions. Langmuir 4, 855–861. Gíslason, S.R., Wolff-Boenisch, D., Stefansson, A., Oelkers, E.H., Gunnlaugsson, E., Sigurdardottir, H., Sigfusson, B., Broecker, W.S., Matter, J.M., Stute, M., Axelsson, G., Fridriksson, T., 2010. Mineral sequestration of carbon dioxide in basalt: a pre-injection overview of the CarbFix project. Int. J. Greenhouse Gas Control 4, 537–545. Golubev, S.V., Pokrovsky, O.S., 2006. Effect of pH and organic ligands on diopside dissolution kinetics. Chem. Geol. 235, 377–389. Golubev, S.V., Bauer, A., Pokrovsky, O.S., 2006. Effect of pH and organic ligands on the kinetics of smectite dissolution at 25 °C. Geochim. Cosmochim. Acta 70, 4436–4451. Govett, G.J.S., 1961. Critical factors in the colorimetric determination of silica. Geochim. Cosmochim. Acta 25, 69–80. Grandstaff, D.E., 1978. Changes in surface area and morphology and the mechanism of forsterite dissolution. Geochim. Cosmochim. Acta 42, 1899–1901.

Grandstaff, D.E., 1986. The dissolution rate of forsteritic olivine from Hawaiian beach sand. In: Colman, S.M., Dethier, D.P. (Eds.), Rates of Chemical Weathering of Rocks and Minerals. Academic Press, pp. 41–57. Guyot, F., Daval, D., Dupraz, S., Martinez, I., Menez, B., Sissmann, O., 2011. CO2 geological storage: the environmental mineralogy perspective. C. R. Geosci. 343, 246–259. Hänchen, M., Prigiobbe, V., Storti, G., Seward, T.M., Mazotti, M., 2006. Dissolution kinetics of forsteritic olivine at 90–150 °C including effects of the presence of CO2. Geochim. Cosmochim. Acta 70, 4403–4416. Hodson, M.E., 2006. Does reactive surface area depend on grain size? Results from pH 3, 25 °C far-from-equilibrium flow-through dissolution experiments on anorthite and biotite. Geochim. Cosmochim. Acta 70, 1655–1667. Huang, W.H., Kiang, W.C., 1972. Laboratory dissolution of plagioclase feldspar in water and organic acids at room temperature. Am. Mineral. 57, 1848– 1859. Innskeep, W.P., Bloom, P.R., 1986. Kinetics of calcite precipitation in the presence of water-soluble organic ligands. Soil Sci. Soc. Am. J. 50, 1167–1172. Jonckbloedt, R.C.L., 1998. Olivine dissolution in sulfuric acid at elevated temperatures, implications for the olivine process, and alternative waste acid neutralizing process. J. Geochem. Explor. 62, 337–346. Katz, J.L., Reick, M.R., Herzog, E.R., Parsiegla, K.L., 1993. Calcite growth inhibition by iron. Langmuir 9, 1423–1430. Kelemen, P.B., Matter, J., 2008. In situ carbonation of peridotite for CO2 storage. Proc. Natl. Acad. Sci. USA 105, 17295–17300. King, H.E., Plumper, O., Putnis, A., 2010a. Effect of secondary phase formation of the carbonation of olivine. Environ. Sci. Technol. 44, 6503–6509. King, H.E., Stimpfl, M., Deymier, P., Drake, M.J., Catlow, C.R.A., Putnis, A., de Leeuw, N.H., 2010b. Computer simulation of water interaction with low-coordinated forsterite surface sites: implications for the origin of water in the inner solar system. Earth Planet. Sci. Lett. 300, 11–18. Kitano, Y., Hood, D.W., 1965. The influence of organic material on the polymorphic crystallization of calcium carbonate. Geochim. Cosmochim. Acta 29, 29–41. Kohler, S.J., Harouiya, N., Chairat, C., Oelkers, E.H., 2005. Experimental studies of REE fractionation during water–mineral interactions: REE release rates during apatite dissolution from pH 2.8 to 9.2. Chem. Geol. 222, 168–182. Lebrón, I., Suárez, D.L., 1998. Kinetics and mechanisms of precipitation of calcite as affected by pCO2 and organic ligands at 25 °C. Geochim. Cosmochim. Acta 62, 405–416. Luce, R.W., Bartlett, R.W., Parks, G.A., 1972. Dissolution kinetics of magnesium silicates. Geochim. Cosmochim. Acta 36, 35–50. Lüttge, A., 2006. Crystal dissolution kinetics and Gibbs free energy. J. Electron Spectrom. Relat. Phenom. 150, 248–259. Lüttge, A., Bolton, E.W., Lasaga, A.C., 1999. An interferometry study of the dissolution kinetics of anorthite: the role of reactive surface area. Am. J. Sci. 299, 652–678. Manley, E.P., Evans, L.J., 1986. Dissolution of feldspar by low-molecular-weight aliphatic and aromatic acids. Soil Sci. 141, 106–112. Marini, L., 2007. Geological sequestration of carbon dioxide: thermodynamics, kinetics and reaction path modeling. Developments in Geochemistry, vol. 11. Elsevier, New York. Maroto-Valer, M.M., Fauth, D.J., Kuchta, M.E., Zhang, Y., Andresen, J.M., 2005. Activation of magnesium rich minerals as carbonation feedstock materials for CO2 sequestration. Fuel Process. Technol. 86, 1627–1645. Mast, A.M., Drever, J.I., 1987. The effect of oxalate on the dissolution rate of oligoclase and tremolite. Geochim. Cosmochim. Acta 51, 2559–2568. Murphy, W.M., Helgeson, H.C., 1987. Thermodynamic and kinetic constraints on reaction rates among minerals and aqueous solutions. III. Activated complexes and the pH-dependence of the rates of feldspar, pyroxene, wollastonite, and olivine hydrolysis. Geochim. Cosmochim. Acta 51, 3137–3153. Murphy, W.M., Helgeson, H.C., 1989. Thermodynamic and kinetic constraints on reaction rates among minerals and aqueous solutions. IV. Retrevial of rate constants and activation parameters for the hydrolysis of pyroxene, wollastonite, olivine, analusite and quartz. Am. J. Sci. 288, 17–101. Noiriel, C., Luquot, L., Madé, B., Raimbault, L., Gouze, P., van der Lee, J., 2009. Changes in reactive surface area during limestone dissolution: an experimental and modeling study. Chem. Geol. 265, 160–170. O’Connor, W.K., Dahlin, D.C., Nilsen, D.N., Walters, R.P., Turner, P.C., 2000a. Carbon dioxide sequestration by direct carbonation with carbonic acid. In: Proc. 25th Int. Tech. Conf. Coal Utilization & Fuel Systems, Coal Technology Assoc., Clearwater, FL. O’Connor, W.K., Dahlin, D.C., Turner, P.C., Walters, R., 2000b. Carbon dioxide sequestration by ex-situ mineral carbonation. In: Proc. 2nd Ann. Dixy Lee Ray Memorial Symp., Am. Soc. Mech. Eng., Washington, DC. Oelkers, E.H., 2001. An experimental study of forsterite dissolution rate as a function of temperature and aqueous Mg and Si concentration. Chem. Geol. 148, 485–494. Oelkers, E.H., Gíslason, S.R., 2001. The mechanism, rates and consequences of basaltic glass dissolution: I. An experimental study of the dissolution rates of basaltic glass as a function of aqueous Al, Si and oxalic acid concentration at 25 °C and pH = 3 and 11. Geochim. Cosmochim. Acta 65, 3671–3681. Oelkers, E.H., Schott, J., 1998. Does organic acid adsorption affect alkali-feldspar dissolution rates? Chem. Geol. 155, 235–245. Oelkers, E.H., Schott, J., 2005. Geochemical aspects of CO2 sequestration. Chem. Geol. 217, 183–186. Oelkers, E.H., Gislason, E.H., Matter, J., 2008. Mineral carbonation of CO2. Elements 4, 333–337.

J. Declercq et al. / Applied Geochemistry 39 (2013) 69–77 Oelkers, E.H., Golubev, S.V., Chairat, C., Pokrovsky, O.S., Schott, J., 2009. The surface chemistry of multi-oxide silicates. Geochim. Cosmochim. Acta 73, 4617–4634. Olsen, A.A., Rimstidt, D., 2008. Oxalate promoted forsterite dissolution at low pH. Geochim. Cosmochim. Acta 72, 1758–1766. Olsson, J., Bovet, N., Makovicky, E., Bechgaard, K., Balogh, Z., Stipp, S.L.S., 2012. Olivine reactivity with CO2 and H2O on a microscale: implications for carbon sequestration. Geochim. Cosmochim. Acta 77, 86–97. Orlando, A., Borrini, D., Marini, L., 2001. Dissolution and carbonation of a serpentine: inferences from acid attack and high P–T experiments performed in aqueous solutions at variable salinity. Appl. Geochem. 26, 1569–1583. Otsuki, A., Wetzel, R.G., 1972. Coprecipitation of phosphate with carbonates in a marl lake. Limnol. Oceanogr. 17, 735–737. Paquette, J., Vali, H., Mucci, A., 1996. TEM study of Pt–C replicas of calcite overgrowths precipitated from electrolyte solutions. Geochim. Cosmochim. Acta 60, 4689–4699. Perrin, D.D., Dempsey, B., Serjeant, E.P., 1981. pKa Predictions for Organic Acids and Bases. Chapman & Hall, London. Pokrovsky, O.S., Schott, J., 1999. Olivine surface speciation and reactivity in aquatic systems. In: Geochemistry of the Earth’s Surface. Balkema, Rotterdam, pp. 461– 463. Pokrovsky, O.S., Schott, J., 2000a. Forsterite surface composition in aqueous solutions: a combined potentiometric, electrokinetic, and spectroscopic approach. Geochim. Cosmochim. Acta 64, 3299–3312. Pokrovsky, O.S., Schott, J., 2000b. Kinetics and mechanism of forsterite dissolution at 25 °C and pH from 1 to 12. Geochim. Cosmochim. Acta 64, 3313–3325. Pokrovsky, O.S., Dupré, B., Schott, J., 2005. Fe–Al-organic colloids control of trace elements in peat soil solutions: results of ultrafiltration and dialysis. Aquat. Geochem. 11, 241–278. Pokrovsky, O.S., Shirokova, L.S., Bénézeth, P., Schott, J., Golubev, S.V., 2009. Effect of organic ligands and heterotrophic bacteria on wollastonite dissolution kinetics. Am. J. Sci. 309, 731–772. Poulson, S.R., Drever, J.I., Stillings, L.L., 1997. Aqueous Si-Oxalate complexing, oxalate adsorption onto quartz and the effect of oxalate upon quartz dissolution rates. Chem. Geol. 140, 1–7. Prigiobbe, V., Mazzotti, M., 2011. Dissolution of olivine in the presence of oxalate, citrate and CO2 at 90 and 120 °C. Chem. Eng. Sci. 66, 6544–6554. Prigiobbe, V., Costa, G., Baciocchi, R., Hänchen, M., Mazzotti, M., 2009. The effect of CO2 and salinity on olivine dissolution kinetics at 120 °C. Chem. Eng. Sci. 64, 3510–3515. Reddy, M.M., 1977. Crystallization of calcium carbonate in the presence of trace concentrations of phosphorous containing anions. I. Inhibition by phosphate and glycerophosphate ions at pH 8.8 and 25 °C. J. Cryst. Growth 41, 287–295. Reddy, M.M., Wang, K.K., 1980. Crystallization of calcium carbonate in the presence of metal ions. I. Inhibition by magnesium ion at pH 8.8 and 25 °C. J. Cryst. Growth 50, 470–480. Reynolds Jr., R.C., 1978. Polyphenol inhibition of calcite precipitation in Lake Powell. Limnol. Oceanogr. 23, 585–597. Rimstidt, D., Brantley, S.L., Olsen, A.A., 2012. Systematic review of forsterite dissolution rate data. Geochim. Cosmochim. Acta 99, 159–178. Rosso, J.J., Rimstidt, J.D., 2000. A high resolution study of forsterite dissolution rates. Geochim. Cosmochim. Acta 64, 797–811. Saldi, G.D., Daval, D., Morvan, G., Knauss, K.G., 2013. The role of Fe and redox conditions in olivine carbonation rates: an experimental study of the rate

77

limiting reactions at 90 and 150 °C in open and closed systems. Geochim. Cosmochim. Acta 118, 157–183. Sanemasa, I., Yoshida, L., Ozawa, T., 1972. The dissolution of olivine in aqueous solutions of inorganic acids. Bull. Chem. Soc. Jpn. 45, 1741–1746. Schott, J., Pokrovsky, O.S., Oelkers, E.H., 2009. The link between mineral dissolution/ precipitation kinetics and solution chemistry. Rev. Mineral. Geochem. 70, 207– 258. Siever, R., Woodfort, N., 1979. Dissolution kinetics and the weathering of mafic minerals. Geochim. Cosmochim. Acta 43, 717–724. Stumm, W., Wollast, R., 1990. Coordination of chemistry of weathering: kinetics of the surface-controlled dissolution of oxide minerals. Rev. Geophys. 28, 53–69. Suess, E., 1970. Interaction of organic compounds with calcium carbonate – I. Association phenomena and geochemical implications. Geochim. Cosmochim. Acta 34, 157–168. Teir, S., Revitzer, H., Eloneva, S., Fogelholm, C., Zevenhoven, R., 2007. Dissolution of natural serpentine in mineral and organic acids. Int. J. Miner. Process. 83, 36–46. Thorseth, I.H., Furnes, H., Tumyr, O., 1995. Textural and chemical effects of bacterial activity on basaltic glass: an experimental approach. Chem. Geol. 119, 139–160. Ullman, W.J., Welch, S.A., 2002. Organic ligands and feldspar dissolution. In: Hellmann, R., Wood, S.A. (Eds.), Water–Rock Interactions, Ore Deposits, and Environmental Geochemistry: A Tribute to David A. Crerar. Geochem. Soc. Special Pub. 7, pp. 3–35. Ullman, W.J., Kirchman, D.L., Welch, S.A., Vandevivere, P., 1996. Laboratory evidence for microbially mediated silicate mineral dissolution in nature. Chem. Geol. 132, 11–17. Van Herk, J., Pietersen, H.S., Schuiling, R.D., 1989. Neutralization of industrial-waste acids with olivine – the dissolution of forsteric olivine at 40–70 °C. Chem. Geol. 76, 341–352. Wang, F., Gaimmar, D.E., 2013. Forsterite dissolution in saline water at elevated temperature and high CO2 pressure. Environ. Sci. Technol. 47, 168–173. Welch, S.A., Ullman, W.J., 1993. The effect of organic acids on plagioclase dissolution rates and stoichiometry. Geochim. Cosmochim. Acta 57, 2725–2736. Welch, S.A., Ullman, W.J., 1996. Feldspar in acidic and organic solutions: compositional and pH dependence of dissolution rate. Geochim. Cosmochim. Acta 60, 2939–2948. Weres, O., Yee, A., Tsao, L., 1981. Kinetics of silica polymerization. J. Colloid Interface Sci. 84, 379–402. Wogelius, R.A., Walther, J.V., 1991. Olivine dissolution at 25 °C: effects of pH, CO2 and organic acids. Geochim. Cosmochim. Acta 55, 943–954. Wogelius, R.A., Walther, J.V., 1992. Olivine dissolution kinetics at the near-surface conditions. Chem. Geol. 97, 101–112. Wolf, G.H., Chizmeshya, A.V.G., Diefenbacher, J., McKelvy, M.J., 2004. In-situ observation of CO2 sequestration reactions using a novel microreaction system. Environ. Sci. Technol. 38, 932–936. Wolff-Boenisch, D., Gislason, S.R., Oelkers, E.H., 2006. The effect of crystallinity on dissolution rates and CO2 consumption capacity of silicates. Geochim. Cosmochim. Acta 70, 858–870. Wolff-Boenisch, D., Wenau, S., Gislason, S.R., Oelkers, E.H., 2011. Dissolution of basalts and peridotite in seawater, in the presence of ligands, and CO2: implication for CO2 sequestration of carbon dioxide. Geochim. Cosmochim. Acta 75, 5510–5525.