Fluorapatite surface composition in aqueous solution deduced from potentiometric, electrokinetic, and solubility measurements, and spectroscopic observations

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Geochimica et Cosmochimica Acta 71 (2007) 5888–5900 www.elsevier.com/locate/gca

Fluorapatite surface composition in aqueous solution deduced from potentiometric, electrokinetic, and solubility measurements, and spectroscopic observations Claire Chaı¨rat a b

a,b

, Eric H. Oelkers

a,*

, Jacques Schott a, Jean-Eric Lartigue

b

Bioge´ochimie et Ge´ochimie Expe´rimentale LMTG-Universite´ de Toulouse-CNRS-IRD-OMP, 14 av. Edouard Belin 31400 Toulouse, France Laboratoire d’Etude du Comportement a` Long Terme LCLT/SECM/DTCD CEA Valrhoˆ BP17171 30207 Bagnols-sur-Ce`ze Cedex, France Received 13 April 2006; accepted in revised form 20 September 2007; available online 5 October 2007

Abstract The surface chemistry of fluorapatite in aqueous solution was investigated using electrokinetic techniques, potentiometric titrations, solubility measurements, and attenuated total reflection infrared spectroscopy. All methods indicate the formation of Ca/F depleted, P enriched altered layer via exchange reactions between H+ and Ca2+, and OH and F at the fluorapatite (FAP) surface. Observations suggest that this leached layer has a di-calcium phosphate (CaHPO4) composition and that it controls the apparent solubility of FAP. Electrokinetic measurements yield an iso-electric point value of 1 ± 0.5 consistent with a negatively charged FAP surface at pH > 1. In contrast, surface titrations give an apparent pH of point of zero charge of 7.7, consistent with a positively charged surface at pH < 7.7. These differences are shown to stem from proton consumption by both proton exchange and dissolution reactions at the FAP surface. After taking account for these effects, FAP surface charge is shown to be negative to at least pH 4 by surface titration analysis.  2007 Elsevier Ltd. All rights reserved.

1. INTRODUCTION The goal of this study is the improved understanding of apatite surface reactivity. This work is motivated by the role of apatite in natural and industrial processes. Apatite weathering controls the phosphorus availability in many terrestrial and marine environments. Phosphorus is (1) essential to micro-organisms and (2) can aid in the remediation of metal-contaminated soils and waters (Xu et al., 1994; Chen et al., 1997). Reactions at the apatite surface also play a major role in fertilizer production (Becker, 1989), medicine (Aoki, 1991), biomineralisation in bones (Mann et al., 1989; Driessens and Verbeeck, 1990), bio-ceramics (Heich and Wilson, 1993), and dental protection (LeGeros, 1991). Moreover, apatite minerals are currently being considered as both possible additives to

*

Corresponding author. Fax: +33 5 61 33 25 60. E-mail address: [email protected] (E.H. Oelkers).

0016-7037/$ - see front matter  2007 Elsevier Ltd. All rights reserved. doi:10.1016/j.gca.2007.09.026

backfill material (Perrone et al., 2002) or as potential actinide and iodine storage hosts for the confinement of nuclear wastes (e.g. Boyer et al., 1998; Guy et al., 2002; Chaı¨rat et al., 2004; Chaı¨rat et al., 2006). Because of its low solubility, apatite dissolution kinetics are controlled by chemical reactions occurring at the mineral surface (Berner, 1981). Quantifying apatite surface chemistry is thus a prerequisite for describing its dissolution rate within the framework of transition state theory (Eyring, 1935) and the surface coordination model (Stumm et al., 1987; Schott, 1990; Stumm and Wieland, 1990). Unlike for oxide, silicate, and carbonate minerals (e.g., Schott et al., 1981; Chou et al., 1989; Arvidson and Mackenzie, 1999; Pokrovsky and Schott, 1999, 2000, 2001; Pokrovsky et al., 1999, 2005; Oelkers, 2001; Oelkers and Schott, 2001), relatively few studies have focused on the surface chemistry of phosphate minerals (Wu et al., 1991; Perrone et al., 2002; Rodenas et al., 2005; Skartsila and Spanos, 2007). Apatite surfaces are thought to consist of two distinct surface groups BCaOH2 þ and BPO . According to the surface

Fluorapatite surface chemistry in aqueous solution

protonation model of Wu et al. (1991), apatite surface protonation proceeds via the formation of „POH surface groups at 5 < pH < 7; apatite surfaces are thus fully protonated at pH < 5. Apatite dissolution rates, however, increase dramatically with decreasing pH at 2 6 pH 6 6 (Guidry and Mackenzie, 2003; Chaı¨rat et al., 2007). This observation suggests an inconsistency between apatite surface chemistry and its dissolution behavior. Towards an improved understanding of fluorapatite (FAP) surface reactivity, we present a detailed description of the fluorapatite–water interface obtained by combining potentiometric surface titrations, electrokinetic measurements, surface spectroscopy analysis, and measurements of fluorapatite solubility as a function of pH. A companion paper (Chaı¨rat et al., 2007) presents experimentally measured fluorapatite dissolution rates and relates these rates to the fluorapatite surface chemistry. 2. MATERIALS AND METHODS 2.1. Fluorapatite samples A natural pegmatitic fluorapatite from Paraı¨ba, Brazil was used in this study. X-ray diffraction analysis shows that this sample is well crystallized, pure apatite. The chemical composition of this apatite was determined using both a Camebax microprobe SX50 using 12 different scan spots and an ELAN 6000 ICP-MS after alkaline fusion for three replicates. The results of these analyses are shown in Tables 1 and 2. The average composition of this apatite was found to be consistent with Ca10.0(PO4)6.0F1.4OH0.6. The sample was initially crushed with a hammer covered by a plastic sheet then ground with an agate mortar and pestle. The resulting powder was sieved to collect specific size fractions. The <10 lm fraction was used without further treatment for the experiments presented in this study. Its BET surface area was 3.54 ± 0.03 m2/g, as determined by eight point krypton adsorption using a Micrometrics ASAP 2010. Table 1 Electron microprobe analysis of the FAP used in the present study Element

Atoms

F Na Mg Al Si P Cl Ca Mn Fe Sr La Ce OHa

1.41 ± 0.08 0.01 ± 0.01 0 0 0.10 ± 0.02 5.85 ± 0.04 0.01 ± 0.01 10.15 ± 0.10 0.01 ± 0.01 0.01 ± 0.01 0 0 0 0.58 ± 0.08

Results are normalised to 18 atoms. a Hydroxide concentration was calculated by subtracting the measured concentrations of F and Cl from 2, the number of anion sites in the FAP structure.

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Table 2 Average composition of the FAP used in the present study deduced from alkaline fusion analysis Ca

P

F + OH

10.08 ± 0.22

6.00

2.16 ± 0.24

Results are provided as atomic ratios normalized to 6.0 phosphorus atoms per mole of apatite.

2.2. Fourier-transformed infrared spectroscopy attenuated total reflectance (FTIR-ATR) To aid in the interpretation of solubility and titration measurements the infrared spectra of three selected FAP samples were measured. The samples analyzed were (1) fresh FAP, (2) FAP following its dissolution for 2 days in a pH 2 solution, and (3) FAP following its dissolution for 7 days in a pH 12 solution. All samples were rinsed repeatedly with methanol then dried at ambient temperature overnight prior to analysis. Spectra were obtained with a Thermo (Nicolet 5700) FTIR spectrometer using ZnSe reflection elements. The spectrometer is purged with dry air prior to each measurement. The spectra were taken using a 8 cm1 resolution over the 3000–650 cm1 region and a 4 cm1 resolution over the 4000–3000 cm1 region. The enhanced resolution was needed to overcome spectral noise over this latter region. In all cases, spectra were obtained by summing up to 256 individual scans and then subtracting this result from a blank scan obtained without FAP powder present in the spectrometer. 2.3. Electrokinetic measurements Microelectrophoresis can be used to characterize the electric double layer (EDL) of dispersive particles (see Hunter, 1989; Van der Wal et al., 1997; Lyklema et al., 1998 for principles and discussions). This technique was used to determine the FAP zeta potential (f) as a function of pH and ionic strength. These measurements also yield the isoelectric point (pHIEP), which is the pH when f = 0. FAP electrophoretic mobilities were determined using a CAD Instrumentation ‘‘Zetaphoremeter IV’’ Z 4000, microelectrophoremeter following the protocol described by Ge´labert et al. (2004) and Pokrovsky and Schott (2004). The electrophoretic mobility of the particles was converted into zeta potential using the Smoluchowski–Helmholtz equation: f¼

ðelE Þ ð4pgÞ

ð1Þ

where f stands for the zeta potential and e, g, and lE represent the dielectric constant, viscosity, and electrophoretic mobility, respectively, of the solution. Note that lE = VE/ E where VE designates the particle velocity and E refers to the electric field. Experiments were performed in solutions having NaCl concentrations from 0.01 to 0.5 M and 1.2 < pH < 12. pH was adjusted by adding HCl or NaOH to this solution such that the total ionic strength varied somewhat as a function of pH. HCl or NaOH was added just prior to mobility measurement which took no more

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C. Chaı¨rat et al. / Geochimica et Cosmochimica Acta 71 (2007) 5888–5900

than 10 min to perform. At least three replicates were carried out at each condition. 2.4. Surface titrations Acid–base titrations of FAP, aimed at determining surface charge as a function of pH, were performed using a DL70ES Mettler Toledo automatic titrator equipped with a 0.1 mV resolution pH-meter. The titration cell was held at a constant temperature of 25 ± 1 C. Because the solubility of many phosphate bearing phases decreases dramatically with increasing pH, we adopted a similar protocol as used by Perrone et al. (2002) for the titration of FAP surfaces. This method consists of adding a small amount of base to the suspension, starting the titration at alkaline pH, and to then titrate this suspension with an acidic solution. This method makes it possible to obtain a complete titration curve with a single titration run. FAP-aqueous solution suspensions were prepared by adding 0.2 g of FAP to 30 mL of three distinct aqueous solutions so that the solid–fluid ratio was 6.7 g/L. The three solutions contained 0.002 M NaOH and 0.01, 0.1, and 0.5 M NaCl, respectively. Prior to the titration, the suspension was bubbled continuously with N2 and allowed to equilibrate for 2 h. To test if this was sufficient time to equilibrate the solutions, several tests were performed by allowing suspensions to equilibrate for 2, 10, and 24 h. No difference was observed in the results obtained from these different equilibration times. The acidic solution used for the titrations consisted of the original 0.01, 0.1, and 0.5 M NaCl background electrolyte plus 5 · 103 M HCl. The suspensions were continuously stirred and purged with N2 gas during the titrations. After electrode calibration, the titration was started by injecting from 10 to 100 lL of the HCl solution to obtain 20 mV pH steps. At each step, the pH was recorded once steady-state was achieved. Steady-state was assumed once the potential drift was less than 0.2 mV/min. Each complete titration took approximately 6 h to perform. Three replicate titrations were made at each condition; each replicate was identical within uncertainty of the corresponding original titration. An identical procedure was performed to determine proton consumption of corresponding blank (FAP-free) solutions. These blank titrations were performed using solutions obtained from filtering the liquid from the original FAP-aqueous solution suspensions just prior to initiating the titration. The proton consumption of each titration ([H+]c,i), where i identifies the titrated solution or suspension, at each pH step was calculated using:

suspension, ci stands for the activity coefficient of the subscripted aqueous species, calculated using the Davies equation, and Ke represents the water ionization constant. The quantity of protons consumed by reactions occurring at the FAP surface ([H+]s) is then obtained from ½Hþ s ¼

½Hþ c;s  ½Hþ c;b S exp

ð3Þ

[H+]c,s and [H+]s,b correspond to the measured proton consumption obtained using Eq. (2) from the titration of a FAP powder and its corresponding blank solutions, respectively, and Sexp denotes surface area of the solid. The surface area used in this calculation is that of the original FAP powder; no provision was made to account for possible surface area changes during the titration. In total, 12 titration runs were performed, with 30 experimental points for each run. Additional titrations were performed to determine the degree to which elements are released from the FAP surface during the above described acid–base titrations. These additional titrations were performed in an identical manner as the titrations described above with the exception that at regular intervals a sample of the solution was taken, filtered and stored for Ca, P, and F analysis. 2.5. Closed-system FAP dissolution experiments Closed-system experiments were performed to determine the apparent solubility of FAP. These experiments were performed in 500 mL Nalgene bottles at 25 C. Three experiments were performed in solutions having initial pH ranging from 3 to 5 and were initiated by placing 3 g of FAP powder and 500 mL of initial solution into the reactors. A fourth experiment, performed in a solution having an initial pH of 9, was initiated by placing 3.5 g of FAP and 150 mL of initial solution into the reactor. All inlet solutions had a 102 M ionic strength obtained by adding appropriate amounts of reagent grade NaCl, HCl, and/or NaOH to demineralized H2O. After sealing each reactor it was placed on a mechanical shaking table in a 25 ± 1 C temperature regulated room. Reactor fluids were regularly sampled using a syringe; the samples were filtered through a 0.45 lm Millipore Nitrocellulose filter prior to analysis. Outlet pH was measured at 25 C within a few hours of sampling using a Metrohm 744 pH meter coupled to a Metrohm Pt1000/B/2 electrode with a 3 M KCl outer filling solution. NIST buffers (pH 4.001, 6.865, and 9.180 at 25 C) were used to calibrate pH electrodes. 2.6. Aqueous solution analysis

ðC a V a  C b V b Þsuspension  ðC a V a  C b V b Þelectrolyte ½H c;i ¼ V pH ðLogðK e ÞpHÞ 10 10 þ  cHþ cOH þ

ð2Þ where V refers to the solution volume, Ca corresponds to the acid concentration, Va represents the volume of acid added to FAP suspension or background electrolyte, CbVb refers to the quantity of base initially added at the

Calcium concentrations were measured by flame atomic absorption spectroscopy using a Perkin Elmer 5100 spectrometer with a detection limit of 0.3 ppm and a precision of better than 4%. Fluoride and phosphate concentrations were determined with a Dionex High Pressure Liquid Chromatograph (HPLC) after elimination of chloride ions using DIONEX Onguard II AgNO3 cartridges with a detection limit of 0.05 and 0.3 ppm, respectively, and a precision of better than 8%.

Fluorapatite surface chemistry in aqueous solution

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are presented in Fig. 1. The reaction times were chosen to insure that the surfaces were sufficiently altered to observe this alteration by IR-ATR techniques. No evidence of secondary mineral precipitation was observed in either FAP leaching experiment. A small negative feature is apparent at 3600 cm1 due to blank spectra subtraction. After acidic dissolution, the IR-ATR spectrum exhibits a new vibration band around 3500 cm1 suggesting protonation of surface PAOH groups and the subsequent incorporation of protons into the FAP structure (Ishikawa et al., 1989, 2000). The anti-symmetric PO4 stretching band is observed at 1024 cm1 on fresh powder and at 1038 cm1 on acidic dissolved FAP. This apparent peak shift is consistent with a change in the coordination environment of phosphate groups. The anti-symmetric PO4 stretching band of octacalcium phosphate (OCP, Ca8(HPO4)2(PO4)4Æ5H2O) is also located at 1038 cm1 Berry and Baddiel (1967) suggesting calcium depletion of the acid dissolved FAP surface. Similarly, Dorozhkin (1997) observed the appearance of OAH bond peaks, and a shift in the phosphate stretching band position on FAP dissolved in aqueous H3PO4 solutions. Note that the HPO4 stretching bands present in OCP and located at 1130, 1123, and 1105 were not observed on the acid dissolved FAP in this study, but due to their location and the fact that they are relatively weak, they may be hidden in the wide PO4 stretching band. These results suggest that a proton–Ca exchange reaction occurs at FAP surface during its dissolution in acidic solution leading to the

2.7. Thermodynamic calculations The standard state adopted in this study is that of unit activity for pure minerals and H2O at any temperature and pressure. For aqueous species other than H2O, the standard state is unit activity of the species in a hypothetical 1 molal solution referenced to infinite dilution at any temperature and pressure. All thermodynamic calculations, including calculation of aqueous activities and saturation states of reactive solutions, were performed using PHREEQC (Parkhurst, 1998) together with its llnl database (Johnson et al., 2000) after adding thermodynamic properties for hydroxyapatite (HAP) taken from Stumm and Morgan (1996), b-tri-calcium phosphate (b-TCP) taken from Wagman et al. (1982), octa-calcium phosphate (OCP) taken from Stumm and Morgan (1996), di-hydrated di-calcium phosphate (DCPD) taken from Wagman et al. (1982), and anhydrous di-calcium phosphate (DCPA) taken from Wagman et al. (1982). These constants are summarized in Table 3. 3. RESULTS 3.1. Surface spectroscopy The IR-ATR spectra of fresh, unreacted FAP, FAP following its dissolution for 2 days in a pH 2 solution, and FAP following its dissolution for 7 days in a pH 12 solution

Table 3 Logarithms of equilibrium constants (Ks) for the dissolution reaction of various apatite precipitation precursors added to llnl database for thermodynamic calculations performed in the present study Abbreviation HAP OCP b-TCP DCPA DCPD

Reaction 2þ

3



Ca10 ðPO4 Þ6 OH2 ¼ 10Ca þ 6PO4 þ 2OH Ca8 ðHPO4 Þ2 ðPO4 Þ4  5H2 O ¼ 8Ca2þ þ 6PO4 3 þ 2Hþ þ 5H2 O Ca3 ðPO4 Þ2 þ 2Hþ ¼ 3Ca2þ þ 2HPO4 2 CaHPO4 ¼ Ca2þ þ HPO4 2 CaHPO4  2H2 O ¼ Ca2þ þ HPO4 2 þ 2H2 O

a

3800

Log Ks

Reference

114 493.8 7.96 6.7 6.6

Stumm and Morgan (1996) Stumm and Morgan (1996) Wagman et al. (1982) Wagman et al. (1982) Wagman et al. (1982)

b

3600

3400

3200

3000 1200

1100

-1 υ cm

non reacted FAP

1000

900

800

-1 υ cm

pH 2 dissolved FAP acidic treated FAP

pH 7treated dissolved basic FAPFAP

Fig. 1. FTIR-ATR spectra of fresh, unreacted FAP, FAP following its dissolution for 2 days in a pH 2 solution, and following its dissolution for 7 days in a pH 12 solution: (a) spectra over the 3800–3000 cm1 range and (b) spectra over the 1200–800 cm1 range. Note that the negative shoulder observed around 3600 cm1 on fresh FAP is an artifact due to blank subtraction.

C. Chaı¨rat et al. / Geochimica et Cosmochimica Acta 71 (2007) 5888–5900

formation of a protonated Ca-leached layer. FAP dissolved in an alkaline solution provokes a splitting of the phosphate vibration band in two separate peaks. A similar splitting of the phosphate bands have been observed in the spectra of a large number of phosphate minerals including hydroxyapatite, octa-calcium phosphate and brushite (Berry and Baddiel, 1967). This band splitting suggests a change in the coordination environment of some phosphate groups, the others remaining in apatitic configuration.

pH 0

2

4

6

8

10

12

0

NaCl KCl

-10

ζ (mV)

5892

-20

3.2. Electrophoretic measurements -30

Fig. 3. Zeta potential of FAP at 25 C as a function of pH in 0.01 M NaCl and KCl (see caption of Fig. 2).

75

this study, I=0.5M this study, I=0.1M this study, I=0.01M 50

Wu et al., 1991 I=0.1M

25

+

[H ]s µmol/m²

Results of electrophoretic measurements are illustrated in Fig. 2. This figure shows FAP zeta potential (f), determined at 25 C and 0.01, 0.1, and 0.5 mol/L NaCl solutions ± HCl and NaOH as a function of solution pH. FAP exhibits negative zeta potential at pH P 1 in all solutions. The f-potential decreases with increasing pH at pH 6 4 and remains nearly constant at 4 6 pH 6 12. A similar negative zeta potential for hydroxyapatite over this pH range was reported by Rodenas et al. (2005). The absolute value of f decreases with increasing ionic strength due to double layer compression. The pHIEP of FAP estimated from these data is 1 ± 0.5. In contrast, Somasundaran (1968) reported a pHIEP = 5.8 based on streaming potential measurements performed in KCl bearing solutions. In an attempt to determine if this difference is due to the identity of the background electrolyte, as suggested by Bell et al. (1973) and Perrone (1999), we performed further zeta potential measurements in KCl solutions and found a pHIEP identical to that measured in NaCl solutions. The results of this latter measurement are illustrated in Fig. 3. 3.3. Surface titrations Results of surface potentiometric titrations are illustrated in Fig. 4. All titration solutions were undersaturated with respect to potential secondary phases including fluorite. For all conditions studied, the excess surface proton concentration ([H+]s) is positive at acidic conditions, decreases with increasing pH, and becomes negative at basic

pH 0

2

4

6

-25

8

10

4

12

0.5M -10

ζ (mV)

6

8

10

pH

0

0.1M -20

0

0.01M

-30

Fig. 2. Zeta potential of FAP at 25 C as a function of pH in 0.01, 0.1 and 0.5 M NaCl. The data points represent the average of five replicate measurements; the size of the error bars surrounding each point was set to encompass the all replicate measurements at each pH.

Fig. 4. Measured concentration of protons consumed by reactions at the FAP surface, [H+]s, during titrations performed at 25 C in solutions having various aqueous NaCl concentrations. [H+]s values were generated from titrations performed in the present study using Eqs. (2) and (3) to correct for the effects of free proton concentration and consumption of protons by the FAP-powderfree blank solution.

conditions. The point of zero charge, deduced from the intersection of titration curves with the x axis is located at pHPZC = 7.7. This value is similar to that reported in the literature; these values range from 6.8 (Somasundaran, 1968) to 8.2 (Wu et al., 1991). The concentration of Ca, P, and F released during parallel titrations are presented in

Fluorapatite surface chemistry in aqueous solution

Fig. 5 and tabulated in the electronic annex. Note the concentrations shown in Fig. 5 are not in equilibrium with the FAP surface, but correspond to the concentration in solution when the steady-state potential criteria of a potential drift of >0.2 mV/min was attained. The concentrations of Ca, P, and F increase with decreasing pH. The presence of these elements in solution indicates that apatite dissolved during the titrations. Moreover, the stoichiometry of Ca, F, and P differs from that of the dissolving apatite, suggesting that some of these elements were removed from the FAP structure via exchange reactions. Similar metal for proton exchange reactions have been postulated to play a role in the surface chemistry of a number of silicates including the alkali feldspars, volcanic glass, and pyroxenes (Schott et al., 1981; Guy and Schott, 1989; Gautier et al., 1994; Oelkers et al., 1994; Oelkers and Gislason, 2001; Oelkers and Schott, 2001; Gislason and Oelkers, 2003). The quantity of fluoride and calcium removed from the FAP structure via exchange reactions can be determined from the differ-

5893

ence between the total measured mass of these elements in the titration fluids and the corresponding mass released via stoichiometric dissolution. This latter mass is calculated from the mass of P in the titration fluids and the stoichiometry of the dissolving FAP; the release of phosphorus can be considered as a tracer of the FAP dissolution reaction since this element does not participate in exchange reactions (Chaı¨rat et al., 2007). Fig. 6 illustrates the computed quantities of Ca and F released to solution via exchange and dissolution reactions as a function of pH. The quantity of Ca released by exchange reactions increases with decreasing pH consistent with its removal via reactions such as: Ca10 ðPO4 Þ6 F1:4 ðOHÞ0:6 þ 2xHþ ¼ Ca10x H2x ðPO4 Þ6 F1:4 ðOHÞ0:6 þ xCa2þ

ð4Þ

In contrast, the quantity of F released by exchange reactions during titrations decreases with decreasing pH which is consistent with its removal via reactions such as: Ca10 ðPO4 Þ6 F1:4 ðOHÞ0:6 þ yOH ¼ Ca10 ðPO4 Þ6 Fð1:4yÞ ðOHÞ0:6þy þ yF :

I=0.01 M

[i] (μmol/L)

1500

1000

500

0 4

6

8

10

12

pH I=0.1 M

[i] (μmol/L)

1500

It can be seen in Fig. 6 that the net amount of F consumed by exchange reactions is negative and pH < 6 indicating that F exchanged into FAP at these conditions. This further confirms the F/OH exchange reaction which, in accord with reaction (5) would favor replacement of OH by F at acidic conditions. The existence of both dissolution and exchange reactions implies that the protons consumed during FAP titration ([H+]s) do not originate solely from sorption/desorption of protons but from several distinct reactions at the FAP surface. A complete description of the protons consumed during these titrations must, therefore, take account of the following reactions: 3.3.1. Stoichiometric dissolution of the bulk solid This reaction releases Ca, P, and F according to:

1000

Ca10 ðPO4 Þ6 F2 ¼ 10Ca2þ þ 6PO4 3 þ 2F 500

ð6Þ

Liberated Ca, P, and F hydrolyze in solution according to Ca2þ þ H2 O ¼ CaOHþ þ Hþ ;

0



4

6

8

10

12

pH I=0.5 M 1500

[i] (μmol/L)

ð5Þ

þ

0

ð7Þ

F þ H ¼ HF ;

ð8Þ

PO4 3 þ Hþ ¼ HPO4 2 ;

ð9Þ

PO4 3 þ 2Hþ ¼ H2 PO4  ;

ð10Þ

and PO4 3 þ 3Hþ ¼ H3 6PO4 0 :

1000

ð11Þ

Within the investigated pH range of 4 6 pH 6 11, only reactions (9) and (10) contribute significantly to the proton balance. The total quantity of protons consumed or liberated by apatite dissolution, [H+]d, can be computed using:

500

0 4

6

8

10

12

pH Fig. 5. Aqueous concentrations of solutions during the FAP titrations. The black diamonds, white triangles, and grey squares refer to the concentrations of Ca, F, and P, respectively. The data plotted in this figure are provided in electronic annex 1.

½Hþ d ¼

1  ð½HPO4 2  þ 2½H2 PO4  Þ S exp

ð12Þ

where Sexp stands for the FAP surface area exposed to solution during the titration and [i] refers to the aqueous concentration of the ith aqueous species. ½HPO4 2  and ½H2 PO4   are determined from measured aqueous P

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C. Chaı¨rat et al. / Geochimica et Cosmochimica Acta 71 (2007) 5888–5900 100

[F ] µmol/L

75 1000

I=0.01 M

-

2+

[Ca ] µmol/L

1500

500

50 25 0 -25 4

0 4

6

8

10

6

12

8

10

12

10

12

pH

pH

100 75

[F-] μmol/L

I=0.1M

[Ca2+] μmol/L

1500

1000

500

50 25 0

0

4

6

8

10

-25

12

4

6

8

pH

pH 1500

100

I=0.5M

[F-] μmol/L

[Ca2+] μmol/L l/

75 1000

500

50 25 0

0 4

6

8

10

pH

12

-25 4

6

8

10

12

pH

Fig. 6. Total amount of calcium and fluoride liberated during FAP as a function of pH: total element liberated (s), liberated by FAP dissolution ( ), liberated by ion exchange reactions (r).

concentration using PHREEQC. This approach is similar to that of Skartsila and Spanos (2007) who observed that apatite dissolution affects strongly the interpretation of potentiometric titrations. 3.3.2. Stoichiometric exchange of 2H+ for one Ca and of one OH for one F at the apatite surface These surface exchange reactions proceed via reactions (4) and (5), respectively. The quantity of fluoride exchanged was found negligible compared to that of calcium (see Fig. 6), thus the Ca2+/H+ exchange was the only exchange reaction considered in the computation of proton consumption stemming from exchange reactions, [H+]e. The protons consumption due to the Ca2+/H+ exchange can therefore be calculated from:   2 10 ð13Þ  ½Ca2þ   ½P : ½Hþ e ¼ S exp 6

The use of Eq. (13) to estimate the quantity of Ca exchanged from the surface is based on the assumption that dissolved P does not adsorb to the FAP surface. This seems likely to be the case as (1) aqueous P is negatively charged at pH > 2 and (2) the zeta-potential measurements suggest that the FAP surface is negatively charged at pH > 1. Calculated proton consumption via the Ca2+/H+ exchange reaction was found to increase with increasing ionic strength. This suggests that sodium ions present in the background electrolyte compete with H+ for calcium at the mineral surface potentially leading to overestimation of [H+]e in high ionic strength solutions. 3.3.3. H+ adsorption on the apatite surface The amount of protons adsorbed onto the FAP surface ([H+]ads) can be computed from the difference between the proton consumption measured from potentiometric titra-

Fluorapatite surface chemistry in aqueous solution 75

75

0.01M

0.1 M

50

50

[H ]i µmol/m²

25

25

+

[H+]i µmol/m²

5895

0

-25

0

-25 4

5

6

7

4

8

6

8

pH

pH 75

50

25

+

[H ]i µmol/m²

0.5 M

0

-25 4

6

8

pH Fig. 7. Computed quantity of protons consumed by dissolution (black curve), exchange (grey curve) and adsorption (grey diamonds) reactions during the FAP titrations preformed in the present study in 0.01, 0.1, and 0.5 M NaCl solutions. Black diamonds represent the total proton consumption of the FAP surfaces during the titrations. The data shown in this figure are tabulated in electronic annex 2–4.

tions and that consumed by dissolution and exchange reactions according to:

proton consumption by dissolution and exchange reactions during surface titration experiments.

½Hþ ads ¼ ½Hþ s  ½Hþ d  ½Hþ e

3.4. Closed system fluorapatite dissolution experiments

ð14Þ

Combining Eqs. (12)–(14) leads to:  1  ½Hþ s :  ð½HPO4 2  þ 2½H2 PO4  Þ  2 ½Hþ ads ¼ S exp  6 ð15Þ  ð½Ca2þ   ½PÞ : 10 The equations described above were used to determine the quantity of protons consumed by each FAP surface reaction. Because dissolved Ca, P, and F concentrations of titration solutions were obtained at different pH than the acid–base titrations, the former values were fit using EXCEL to quadratic equations to interpolate them to the pH of the latter titrations. The results of proton consumption calculations are shown in Fig. 7 and provided in tables as part of an electronic annex. Proton consumption measured during the FAP titrations are dominated by dissolution and exchange reactions. The quantity of protons consumed by adsorption reactions determined by subtracting total proton consumption by that consumed by dissolution and exchange, in accord with Eqs. (13) and (14), are also shown in this figure. In all cases, the computed quantity of protons consumed by adsorption is negative to at least pH 4. Although it is not possible to perform accurate FAP surface titrations down to pH 1 due to the rapid FAP dissolution rate at strongly acid conditions, this result illustrates that the apparent differences between pHZPC and pHIEP reported for multi-oxide silicates are likely due to

The results of closed-system FAP dissolution experiments at 25 C as a function of pH are listed in Table 4. After a transient time, element concentrations and pH become constant within the experimental uncertainty suggesting attainment of a steady-state between the solution and solid. It should be noted that as was the case during the titrations described above, the initial release of Ca and F is non-stoichiometric with respect to phosphorus. A representative variation of aqueous concentrations and element ratios measured during closed system experiments is shown in Fig. 8. Activities of the aqueous species present in the final equilibrated solutions were used to calculate fluorapatite reactions quotients (QFAP) using QFAP ¼ ðaCa2þ Þ10 ðaPO4 3 Þ6 ðaF Þ2

ð16Þ

where ai is the activity of the subscripted aqueous species. Note that this FAP formula differs somewhat from that deduced from chemical analyses but this difference has negligible effects on the calculations presented below. Resulting steady-state reaction quotients are plotted as a function of pH in Fig. 9. It can be seen that calculated steady-state QFAP values are pH dependent, increasing from 136 to 95 with increasing pH. Note that the final equilibrated solutions were not supersaturated with respect to any potential secondary phase except igneous hydroxyapatite

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C. Chaı¨rat et al. / Geochimica et Cosmochimica Acta 71 (2007) 5888–5900

Table 4 Measured pH and aqueous concentrations during closed-system FAP dissolution experiments at 25 C and the corresponding logarithm of the FAP ion activity product (Q) consistent with reaction (3) Ref.

Elapsed time (days)

pH

[Ca]

[P] (lmol kg1)

[F]

Log(Q)

Experiment B1FAP3 B1FAP3-1 B1FAP3-2 B1FAP3-3 B1FAP3-4 B1FAP3-5 B1FAP3-6 B1FAP3-7 B1FAP3-8 B1FAP3-9 B1FAP3-10 B1FAP3-11 B1FAP3-12 B1FAP3-13

1 4 8 11 18 25 32 36 43 63 71 82 95

3.66 4.42 4.51 4.68 4.70 4.83 4.93 5.03 5.28 5.51 5.48 5.57 5.74

521.5 666.0 678.0 688.8 656.6 673.8 670.4 655.0 651.0 659.9 665.9 676.0 662.1

287.0 350.2 359.1 355.7 364.7 363.3 359.8 364.3 364.2 358.3 363.6 346.8 361.7

89.6 117.4 123.2 119.2 124.0 121.2 127.9 124.7 121.4 121.6 123.1 122.3 126.5

138.2 127.0 125.8 123.7 123.5 121.9 120.7 119.6 116.7 113.9 114.2 113.2 111.2

Experiment B2FAP4 B2FAP4-1 B2FAP4-2 B2FAP4-3 B2FAP4-4 B2FAP4-5 B2FAP4-6 B2FAP4-7 B2FAP4-8 B2FAP4-9 B2FAP4-10 B2FAP4-11 B2FAP4-12 B2FAP4-13

1 4 8 11 18 25 32 36 43 63 71 82 95

5.52 6.07 6.27 6.30 6.30 6.41 6.42 6.45 6.69 6.61 6.44 6.75 6.76

80.1 88.8 92.2 92.4 94.9 95.6 96.5 99.5 102.0 106.8 109.2 113.7 115.1

31.9 34.9 33.9 35.5 37.0 36.7 36.1 37.4 37.4 38.5 38.3 40.3 40.3

17.1 12.6 12.6 12.1 12.7 14.8 13.2 15.8 13.7 14.8 16.3 16.0 19.1

130.8 124.0 121.7 121.2 121.0 119.6 119.6 118.9 116.4 116.9 118.5 115.0 114.7

Experiment B3FAP5 B3FAP5-1 B3FAP5-2 B3FAP5-3 B3FAP5-4 B3FAP5-5 B3FAP5-6 B3FAP5-7 B3FAP5-8 B3FAP5-9 B3FAP5-10 B3FAP5-11 B3FAP5-12 B3FAP5-13

1 4 8 11 18 25 32 36 43 63 71 82 95

6.25 6.46 6.77 6.62 6.68 6.77 6.74 6.74 6.96 6.79 6.60 6.78 6.82

32.8 45.5 49.8 49.7 55.2 56.1 58.0 61.9 63.0 71.0 71.5 70.2 74.5

6.5 9.8 11.9 12.6 13.3 14.8 15.2 15.1 16.1 17.6 18.1 17.5 19.5

1.1 4.7 20.3 1.2 6.3 6.4 6.2 8.3 8.9 9.5 8.9 1.7 9.5

132.7 126.7 121.3 125.1 122.5 121.2 121.3 120.8 118.4 119.2 121.1 120.9 118.4

Experiment B0FAP9 BO_F-1 BO_F-2 BO_F-3 BO_F-4 BO_F-5 BO_F-6 BO_F-7 BO_F-8 BO_F-9

8 14 25 68 133 140 160 188 207

9.80 — — 9.90 9.63 9.70 9.84 9.68 9.23

124.0 148.2 160.1 150.6 112.7 126.6 120.7 139.5 131.2

29.8 46.1 52.8 — 6.3 40.1 40.4 61.8 34.2

18.6 22.3 23.4 27.9 25.1 38.9 37.1 47.9 28.9

95.4 93.6 93.0 92.7 100.3 94.4 94.1 93.0 97.0

and b-TCP, which do not form at low temperature (Viellard and Tardy, 1984). These results suggest that the composition of these final equilibrated solutions may be controlled by a surface layer having a composition distinct from that of FAP.

Apatite precipitation is known to be preceded by the formation of various calcium phosphate phases which may ultimately transform into apatite. Brown (1966) suggested that precipitation of octa-calcium phosphate (OCP) (Ca8(HPO4)2(PO4)4ÆnH2O) precedes apatite formation,

Fluorapatite surface chemistry in aqueous solution

were found to vary significantly over the investigated pH range. In contrast, reaction quotient for the di-calcium phosphate dissolution reaction given by:

10

1.5

5897

CaHPO4  nH2 O ¼ Ca2þ þ HPO4 2 þ nH2 O pH

6

4

[i]x10 mol/kg

8 1.0

0.5

ð19Þ

are found to be close to constant over the investigated pH range. The logarithm of the reaction quotient found for this reaction is 9.6 ± 0.6.

4

4. DISCUSSION 0.0 0

25

50

75

2 100

elapsed time (days) 5

Atomic ratio

4

3

2

1

0

0

25

50

75

100

elapsed time (days) Fig. 8. Fluid evolution during closed-system FAP dissolution experiment B2FAP4. (a) Reactive solution concentration as a function of time. Black diamonds, grey squares, white triangles, and stars correspond to the calcium concentration, phosphorus concentration, fluoride concentration, and pH, respectively. (b) Reactive fluid concentration ratios. Black diamonds and grey squares correspond to measured Ca/P and F/P concentration ratios in the reactive solution whereas the dashed black and grey lines represents these ratios in the dissolving FAP.

Destainville et al. (2003) proposed that apatitic tri-calcium phosphate (apatitic TCP) (Ca9HPO4(PO4)5F) is a precursor to apatite formation, Francis and Web (1971) and Neuman and Bareham (1975) suggested that di-calcium phosphate (DCP) (CaHPO4ÆnH2O, where n = 0 or 2) is a precursor to apatite formation, and Eanes et al. (1966), Brecevic and Furedi-Milhofer (1972), and Termine et al. (1970) proposed that amorphous calcium phosphate (ACP) is an apatite precursor. Dorozhkin (1997) suggested that these precursors may also be intermediaries during the FAP dissolution process. To test if one of these similar minerals is controlling the apparent solubility of FAP, activities of all aqueous species present in the final equilibrated solutions were used to calculate the reaction quotient of these precursor phases. Results are presented in Fig. 9 and Table 5. Reaction quotients for the apatitic tri-calcium phosphate and octa-calcium phosphate dissolution reactions given by: Ca9 ðHPO4 ÞðPO4 Þ5 F ¼ 9Ca2þ þ HPO4 2 þ 5PO4 3 þ F ð17Þ Ca8 ðHPO4 Þ2 ðPO4 Þ4 ¼ 8Ca2þ þ 2Hþ þ 6PO4 3

ð18Þ

Results of ATR spectroscopy as well as closed-system FAP dissolution experiments suggest that a Ca/F depleted layer is formed on the FAP surface upon its introduction into aqueous solution. This Ca and F depleted layer could be formed by the exchange of two hydrogen for one calcium atom, and one hydroxyl for one fluoride at the FAP surface or by a coupled Ca and F removal and phosphate hydrolysis. In addition, the ATR data also suggests that some protons may have penetrated into the FAP structure. Proton penetration into the surface lattice of FAP under acidic conditions has already been reported by Dorozhkin (1997) using diffuse reflective infrared Fourier transform spectroscopy. Comparison of zeta potential measurements, where only the first layer determines the electrokinetic properties and the pHIEP of the mineral/water interface, and surface titration data indicates that several reactions are involved in the uptake of H+ by FAP surfaces. This leads to an apparent discrepancy between pHZPC and pHIEP where the former is apparently substantially higher than the latter. Similar apparent discrepancies between pHZPC and pHIEP due to protons uptake in the surface lattice was already observed for silicate minerals (Brantley and Stillings, 1996, 1997; Pokrovsky and Schott, 2000). It seems likely that this discrepancy originates from failure to accurately account for the consumption of protons by dissolution and exchange reactions during the titration reactions. Previous surface speciation models for fluorapatite (Wu et al., 1991) and carbonate apatite (Perrone, 1999) have been based solely on potentiometric titration data. It was assumed in these previous studies that all protons consumed during titration were due to adsorption onto surface sites. In contrast, the present study demonstrates that such an interpretation is flawed. For example, Wu et al. (1991) assumed that proton consumption during potentiometric titration of FAP surfaces at 5 < pH 6 8 reflects the protonation of surface phosphate groups according to: BPO þ Hþ ¼ BPOH:

ð20Þ

This interpretation is, however, inconsistent with (1) zeta potential measurements which show that FAP surface is negatively charged to pH P 1, (2) negative net proton adsorption in potentiometric titrations to at least pH 4, and (3) the analogy generally observed between surface and aqueous hydrolysis constants (Schlinder and Stumm, 1987). The logarithm of the first phosphoric acid hydrolysis constant is 2.2 at 25 C, which differs dramatically with the corresponding logarithm for the first protonation constant of phosphate groups at the apatite surface reported by

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C. Chaı¨rat et al. / Geochimica et Cosmochimica Acta 71 (2007) 5888–5900 -60

log(QFAP) = 6.05 pH- 155

log Q = 5.391 pH- 137

R2 = 0.90

2

R = 0.90

-80

Log (Q)

Log (QFAP)

-80 -100 -120 -140

-110

Ca9(HPO4)(PO4)5F

FAP

-160

-140 2

4

6

8

10

2

12

4

6

pH -60

12

log Q= 0.196 pH - 10.66 2 R = 0.43

2

Log (Q)

R = 0.90

Log (Q)

10

-6

log Q = 4.74 pH - 118

-90

Ca8 (HPO4)2(PO4)4 -120

8

pH

-9

CaHP O 4

-12 2

2

4

6

8

10

4

6

8

10

12

12

pH

pH

Fig. 9. Variation of the logarithm of reaction quotients of various phosphate bearing minerals computed from the final steady-state solution concentration of closed-system FAP dissolution experiments performed in this study ( ) and by Harouyia et al. (2007) (d). The identity of each mineral and the result of a linear regression of these data are provided in the figure.

Table 5 Saturation index (Log(X)) of potential secondary products at steady-state of FAP closed-system dissolution experiments

B1FAP3 B2FAP4 B3FAP5 B0FAP9

HAP

Fluorite

Portlandite

b-TCP

OCP

DCPA

DCPD

8.0 7.3 9.9 19.7

1.2 3.8 4.6 3.0

14.7 13.4 13.4 7.4

1.3 1.5 2.4 5.5

15.8 17.8 20.4 2.7

1.8 2.6 3.1 2.1

1.9 2.7 3.2 2.2

Saturation indexes were calculated with PHREEQC together with llnl database according to Log(X) = Log(Q/Ks) where Q designates the ionic activity product and Ks the solubility product of the solid.

Wu et al. (1991) of 6.6. It follows that the Wu et al. (1991) FAP surface speciation model, and similar models based on the results of the acid–base titration of multi-component solids need to be reconsidered by taking account of all proton consuming reactions at the mineral–solution interface. FAP reaction quotients computed from closed-system experiments performed in this study and by Harouiya et al. (2007) increase from 136 to 95 with pH increasing from 3.8 to 9.2. Similarly, Levinskas and Neuman (1955) concluded that a single solubility product expression could not describe FAP solubility. Nevertheless, the same steady-state solution composition data yields di-calcium phosphate reaction quotients that are close to pH independent. This observation suggests that a leached layer of this composition may control the apparent solubility of FAP. The retrieved di-calcium phosphate reaction quotient is lower than the di-calcium phosphate equilibrium constant reported by Wagman et al. (1982) and Valsami-Jones (2001), suggesting that the leached layer formed on the dissolving FAP surface is slightly more stable that pure DCP. Reaction quotients computed assuming the leached layer has different stoichiometry

than that of di-calcium phosphate yielded values that varied with pH. Moreover, the formation of a leached layer on the FAP surface with a di-calcium phosphate composition is consistent with both the spectroscopic analyses and solution composition data presented above. 5. CONCLUDING REMARKS The surface chemistry of fluorapatite in aqueous solution was investigated using electrokinetic techniques, potentiometric titrations, closed-system dissolution experiments, and FTIR-ATR spectroscopy. All methods indicate the formation of a Ca/F depleted, P enriched altered layer via exchange reactions between H+ and Ca2, and OH and F at FAP surface. As a result, the FAP surface composition in contact with aqueous solution is always different from its bulk composition. Both dissolution and the exchange reactions affect strongly the H+ consumption observed during surface titrations experiments. These results combined with electrophoretic measurements, which indicate that FAP surface is negatively charged at pH P 1, suggest that pro-

Fluorapatite surface chemistry in aqueous solution

tonation of apatite „PO surface groups is negligible at pH P 3. Although the results presented in this study cannot unambiguously identify the exact stoichiometry of the leached layer formed at FAP surface, evidence suggests it is similar to di-calcium phosphate. This evidence includes (i) spectroscopic observations, (ii) elementary analysis, (iii) the Ca and F impoverishment and proton enrichment and (iv) reactions quotient of apatite determined from closedsystem dissolution experiments. This phase has also been previously proposed as a precursor of precipitated apatite (Francis and Web, 1971; Neuman and Bareham, 1975). The FAP surface chemistry deduced in this study will be used together with measured FAP dissolution rates as a function of pH and solution composition to quantify the FAP dissolution mechanism in Chaı¨rat et al. (2007). APPENDIX A. SUPPLEMENTARY DATA Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.gca. 2007.09.026. REFERENCES Aoki H. (1991) Science and Medical Application of Hydroxyapatite. Japanese Association of Apatite Science, Tokyo, p. 268. Arvidson R. S. and Mackenzie F. T. (1999) The dolomite problem: control of precipitation kinetics by temperature and saturation state. Am. J. Sci. 299, 257–288. Becker P. (1989) Phosphates and Phosphoric Acid, Fertiliser Science and Technology Series, second ed. Dekker, New York, p. 740. Bell L. C., Posner A. M. and Quirck J. P. (1973) The point of zero charge of hydroxyapatite and fluorapatite in aqueous solutions. J. Colloid Interface Sci. 42, 250–261. Berner R. A. (1981) Kinetics of weathering and diagenesis. Rev. Min. 8, 111–134. Berry E. E. and Baddiel C. B. (1967) Some assignments in the infrared spectrum of octacalcium phosphate. Spectrochim. Acta A 23, 1781–1792. Boyer L., Savariault J.-M., Carpena J. and Lacout J.-L. (1998) Neodymium-substituted britholite compound. Acta Cryst. C 54, 1057–1059. Brantley S. L. and Stillings L. (1996) Feldspar dissolution at 25 C and low pH. Am. J. Sci. 296, 101–127. Brantley S. L. and Stillings L. (1997) Feldspar dissolution at 25 C and low pH-reply. Am. J. Sci. 297, 1021–1032. Brecevic L. J. and Furedi-Milhofer H. (1972) Precipitation of calcium phosphates from electrolyte solutions II: the formation and transformation of precipitates. Calc. Tissue Res. 10, 82–90. Brown W. E. (1966) Crystal growth of bone mineral. Clin. Orthopaed. 44, 205–220. Chaı¨rat C., Oelkers E. H., Ko¨hler S. and Harouiya N. (2004) An experimental study of the dissolution rates of apatite and britholite as a function of solution composition and pH from 1 to 12. In Water–Rock Interaction (R. B. Wanty and R. R. Seal, co-editors of the Proceedings) vol. 671–674. Chaı¨rat C., Oelkers E. H., Schott J. and Lartigue J.-E. (2006) An experimental study of the dissolution rates of Nd–britholite, an apatite-structured actinide-bearing waste storage host analogue. J. Nucl. Mater. 354, 14–27. Chaı¨rat C., Schott J., Oelkers E. H., Lartigue J.-E. and Harouyia N. (2007) Kinetic and mechanism of fluorapatite dissolution at

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