Calcium Phosphate Overgrowth on Silicate Sand - ACS Publications

CRYSTAL GROWTH & DESIGN

Calcium Phosphate Overgrowth on Silicate Sand Iosif T. Hafez,†,‡ Christakis A. Paraskeva,*,†,‡ Asimina Toliza,‡ Pavlos G. Klepetsanis,†,§ Petros G. Koutsoukos,†,‡ Øyvind Gustavsen,# Terje Østvold,# and Alkiviades C. Payatakes†,‡ Institute of Chemical Engineering and High-Temperature Processes-FORTH, GR 26 500, Patras, Greece, Chemical Engineering Department, UniVersity of Patras, GR 26 500, Greece, Pharmacy Department, UniVersity of Patras, GR 26 500, Greece, and Department of Chemistry, Norwegian UniVersity of Science and Technology, 7491 Trondheim, Norway

2006 VOL. 6, NO. 3 675-683

ReceiVed August 10, 2005; ReVised Manuscript ReceiVed NoVember 22, 2005

ABSTRACT: Loose sand formations greatly reduce the efficiency of oil pumping from sandy soil reservoirs. Studies of the in-situ deposition of calcium phosphate salts gave promising results, which indicated that consolidation may be achieved through the formation of salt bridges between sand grains. The formation of micron-sized platelike crystallites is desirable for the achievement of maximum sand consolidation with minimal loss of fluid permeability. In the calcium phosphate system, this crystal morphology corresponds to the formation of octacalcium phosphate (Ca8H2(PO4)6‚5H2O, OCP). In the present work, precipitation of calcium phosphate phases from supersaturated solutions on silicate sand substrates was investigated. Because OCP is thermodynamically unstable, the conditions at which this phase may be stabilized was studied with kinetics measurements in batch reactors. Experiments were done in supersaturated solutions seeded with silicate sand over the temperature range from 25 to 70 °C and initial solution pH values 6.0, 7.0, and 8.0. The ionic strength adjusted with sodium chloride was 0.1 and 0.5 M; the solution supersaturation was varied using different calcium and phosphate concentrations (2-20 mM). Clean, dry sand with an average grain size diameter < 0.2 mm was used to seed supersaturated calcium phosphate solutions to determine the influence of the sand grains on the kinetics of formation of the precipitating phase and to investigate the extent of stabilization of calcium phosphates less stable than the thermodynamically most stable hydroxyapatite (Ca5 (PO4)3OH, HAP). A detailed parametric study showed that the experimental conditions of the supersaturated solutions determined the nature of the calcium phosphates forming. Dicalcium phosphate dihydrate (CaHPO4‚2H2O, DCPD) the least stable thermodynamically crystalline form of calcium phosphate formed predominantly at 25 °C and initial pH values lower than 7.0. In the presence of sand grain substrates, DCPD was stabilized to some extent even at relatively high pH. At 50 °C, the formation of DCPD was found to result in a drop of the solution pH to 6.0. At 70 °C, however, no DCPD was detected over the pH range investigated. OCP and HAP were the predominant phases at 50 and 70 °C both in the presence and in the absence of sand. Sand was found to favor the formation of OCP at higher pH values. The ionic strength of the supersaturated solutions adjusted with sodium chloride did not have any appreciable effect on the nature of the calcium phosphate phases nucleated on the sand grains. Introduction Oil production efficiency is seriously affected by the fact that large amounts of solids are pumped from oil reservoirs in loose sandy formations. A number of remediation methods and techniques have been developed for the amelioration of the production process, including in situ chemical consolidation with various agents1-4 and the use of mechanical steel screens and gravel packs between the sandy formation and the screen.5 The chemical consolidation methods involve expensive toxic chemicals posing threats to the environment and the neighboring ecosystems. It has been reported that satisfactory sand bed consolidation may be achieved with in situ deposition of sparingly soluble salts such as calcium phosphate from the respective supersaturated solutions.6,7 The main advantage of this environmentally friendly approach is that it is easily applied, does not involve expensive chemicals, and is reversible. The principle of the consolidation method is the formation of salt bridges between the sand grains, which depending on the nature and the morphology of the crystals deposited, may provide for controlled oil permeability of the consolidated sand formations. The formation of calcium phosphate crystals bridges between sand grains is a promising system as it involves crystals that * To whom correspondence should be addressed. E-mail: takisp@ chemeng.upatras.gr. † Institute of Chemical Engineering and High-Temperature ProcessesFORTH. ‡ Chemical Engineering Department, University of Patras. § Pharmacy Department, University of Patras. # Norwegian University of Science and Technology.

may bind strongly onto sand grains and can be easily dissolved by washing with dilute mineral acid media (pH 4.0 or less is sufficiently effective). The solution chemistry of calcium phosphates however is rather complex, and a number of crystalline phases with widely different physicochemical characteristics (e.g., morphology and solubility) may be formed from the respective supersaturated solutions. As a result, the various phases, which may be formed by nucleation and crystal growth in supersaturated solutions, are expected to exhibit different sand consolidation potentials. The calcium phosphate phases at neutral pH are in the order of decreasing solubility: dicalcium phosphate dihydrate (DCPD, CaHPO4‚2H2O), dicalcium phosphate anhydrous (CaHPO4, DCPA), β-tricalcium phosphate (Ca3(PO4)2, β-TCP) octacalcium phosphate (OCP, Ca8H2(PO4)6‚ 5H2O) and hydroxyapatite (HAP, Ca5 (PO4)3OH). The formation of calcium phosphate phases depends on pH, temperature, ionic strength, and level of supersaturation, and the relative stability of the various phases may change depending on these parameters.8 It has earlier been reported that the formation of OCP and HAP crystallites resulted in a satisfactory consolidation of sand formations.6,7 The consolidation efficiency was attributed to the formation of the crystallites on the surface of the sand grains resulting in the formation of salt bridges. The formation of OCP crystallites was found to yield a more efficient bridging in comparison with HAP. OCP however is thermodynamically unstable, hydrolyzing in the presence of water into HAP.9-12 Kinetic stabilization of OCP at ambient temperature has been reported for very narrow pH range strongly depending on

10.1021/cg050403g CCC: $33.50 © 2006 American Chemical Society Published on Web 01/21/2006

676 Crystal Growth & Design, Vol. 6, No. 3, 2006

temperature.13-16 The presence of silicate substrates is expected to modify the precipitation conditions and subsequently the nature of the calcium phosphate phases forming and the stability of unstable transient phases.17-20 Little is however known concerning the kinetics of the formations of calcium phosphates at conditions similar to those prevalent in oilfield formations, i.e., as a function of the solution temperature and the presence of silicic substrates. Knowledge of these parameters is necessary for the optimization of the sand layers consolidation process. In the present work, which is a continuation of our first report on the possibility to achieve consolidation of sand beds, we report the influence of important parameters such as temperature, pH, salinity, ionic strength, initial calcium, and phosphate concentrations and of the presence of sand on the formation of the calcium phosphate phases from supersaturated solutions. Solutions supersaturated with respect to all calcium phosphates both in the absence and in the presence of silicic sands have been used. The investigation is focused on the identification of the composition of the mineral phases precipitating from the supersaturated solutions over a wide temperature range 20-70 °C. The experimental conditions most favorable for the stabilization of OCP were determined. Experimental Section All experiments were done in a 600 mL double walled vessel thermostated by water circulating through a constant-temperature bath. Four series of experiments were performed at 20, 25, 50, and 70 °C, respectively. The stock solutions were prepared from crystalline, reagent-grade chemicals (Merck, pro analisi), using triply distilled water. Next, the solutions prepared were filtered through membrane filters (0.2 µm Millipore). Calcium chloride and potassium dihydrogen phosphate stock solutions were standardized by atomic absorption spectroscopy and phosphovanadomolybdate spectrophotometric methods, respectively. The sodium chloride stock solution was prepared from crystalline sodium chloride (dried at 105 °C overnight). The supersaturated solutions were prepared in the reactor by simultaneous, rapid mixing of equal volumes, 250 mL each, of two calcium chloride and potassium dihydrogen phosphate solutions (stoichiometry of total calcium, Cat: total phosphate, Pt ) 1:1). The ionic strength of working solutions was adjusted by the addition of the appropriate volume of sodium chloride stock solution. In this study, the experiments at 25, 50, and 70 °C were done at ionic strength values equal to 0.1 and 0.5 M. Next, the supersaturated solutions were degassed for at least 2 h with water vapor saturated nitrogen. Past the lapse of the degassing time, the solution pH was adjusted to the appropriate value by the addition of 0.1 M standard sodium hydroxide solution (Merck, Titrisol). The initial pH values tested in this work were 6.0, 7.0, and 8.0, respectively. The solution pH was measured by a combination glass||Ag/AgCl electrode (Ingold), standardized, before and after each experiment with NBS standard buffer solutions (potassium hydrogen phthalate and Borax).21 The Ca2+ ion activity was measured by a Ca2+ ion selective electrode (ISE) calibrated using standard calcium solutions in the presence of phosphate and at the same ionic strength as the solutions in which measurements were done. The calibration was done at conditions suitable to avoid spontaneous precipitation. Starting from the more dilute to the more concentrated solutions, it was verified that the sensitivity of the electrode was adequate to monitor changes of the calcium ion activity during the course of precipitation. The working solutions were stirred with a magnetic stirrer with a Teflon-coated stirring bar. After pH adjustments, the pH of the working solutions as a function of the reaction time was monitored with a data acquisition system (PCL-818L, Advantech) connected with a personal computer. The solution pH and Ca2+ ion activities were monitored for the experiments done at 20 °C. Upon the start of precipitation, the solution pH decreased as a result of the solid formation. Monitoring of the solution pH proceeded until constant solution pH. Additional series of experiments were done in which the supersaturated solutions were seeded with 5 g of silicate sand, with an average grain diameter < 0.2 mm.

Hafez et al. At the end of the experiments at 25, 50, and 70 °C, the solutions were filtered through membrane filters (Millipore 0.22µm) and the collected solid was dried at 70 °C overnight. The dried solids were characterized by powder XRD (Philips, PW 1830/40, Cu KR radiation), and the morphology of the crystals was examined by scanning electron microscopy (SEM) (JEOL, JSM-5200) equipped with an EDX microanalysis probe (Link, Oxford). In all experiments, the pH profiles as a function of time showed an inflection point. It was thus decided to investigate further this point by monitoring the free calcium ion activity during the course of precipitation. Samples of the solid precipitate were taken before the pH change, which occurred about 40-50 min past the initiation of the experiments done at 20 °C and also at the end of the precipitation process.

Results and Discussion The mean per ion driving force for the formation of the calcium phosphate phases is the change in Gibbs free energy, ∆G, for transfer from the supersaturated to equilibrium state and is

∆G ) -

()

RT IP ln 0 ν Ks

(1)

where R is the gas constant, T is the absolute temperature, IP is the ion activity product of the precipitating phase, and ν (m + n) is the number of ions in the respective phase AmBn. For DCPD and DCPA ν is 2, for TCP ν ) 5, for OCP ν ) 8, and for HAP ν ) 9. In eq 1, K0s is the molar thermodynamic solubility product of the phase considered. The supersaturation ratio, SR, is defined according to eq 2:

SR )

IP K0s

(2)

The expressions for the ionic activity products, for each calcium phosphate phase, that is formed at examined temperature range (25-70 °C) are given by

DCPD and DCPA

IP ) [Ca2+][HPO42-]γ22

(3)

TCP

IP ) [Ca2+]3[PO43-]2γ32 γ23

(4)

OCP

IP ) [Ca2+]4[H+][PO43-]3γ1γ42 γ33 (5)

HAP

IP ) [Ca2+]5[OH-][PO43-]3γ1γ52 γ33

(6)

where γ1 is the activity coefficient for the monovalent, γ2 is for the divalent, and γ3 is for the trivalent ions. The computation of the activity coefficients and of the analytical ion concentrations needed for the calculation of the solution supersaturation with respect to the various calcium phosphate phases precipitating was done by using the computer code HYDRAQL.22 In this code, the speciation problem is solved by successive approximations with an iterative process for the ionic strength, I.23 The system of equations solved includes the mass balance equations for total calcium and total phosphate, Pt, the initial solution pH, and the electroneutrality condition. The activity coefficients of z-valent species, γz, were calculated by the extended DebyeHu¨ckel equation proposed by Davies.24,25

{

-log γz ) Az2

I1/2 - 0.3I I1/2 + 1

}

(7)

The Davies equation was used for calculations in solutions with ionic strength up to 0.5 M. In eq 7, A is the Debye-Hu¨ckel constant (0.5115 at 25 °C).

Calcium Phosphate Overgrowth on Silicate Sand

Crystal Growth & Design, Vol. 6, No. 3, 2006 677

Table 1. Equilibrium Constants for the Calcium Phosphate System at Different Temperaturesa equilibrium H3PO4 T H+ + H2PO4 H2PO4- T H+ + HPO24 HPO42- T H+ + PO34 H+ + OH- T H2O 0 Ca2+ + HPO24 T CaHPO4 + Ca2 + + H2POT CaH PO 2 4 4 Ca2+ + PO3T CaPO 4 4 Ca2 + + OH- T CaOH+ CaHPO4 (DCPA) T Ca2+ + HPO24 CaHPO4‚2H2O(DCPD) T Ca2++ HPO24 + Ca4H(PO4)3‚2.5H2O(OCP) T 4Ca2++ 3PO34 +H Ca3(PO4)2(TCP) T 3Ca2++ 2PO34 Ca5(OH)(PO4)3(HAP) T 5Ca2++ 3PO34 +OH a

K25°C 10-3

7.11 × 6.30 × 10-8 4.73 × 10-13 1.004 × 10-14 398 27.54 2.95 × 106 16.98 2.04 × 10-7 1.26 × 10-7 2.51 × 10-50 1.19 × 10-29 3.83 × 10-59

K50°C 10-3

5.79 × 6.71 × 10-8 7.42 × 10-13 6.66 × 10-14 645 46.28 5.09 × 106 22.62 1.62 × 10-7 9.2 × 10-8 3.78 × 10-51 1.90 × 10-31 1.54 × 10-59

K70°C 5.86 × 10-3 6.00 × 10-8 8.82 × 10-13 1.67 × 10-13 944 69.98 7.94 × 106 28.51 1.18 × 10-7 1.00 × 10-51 3.89 × 10-34 1.41 × 10-59

From Ref 26.

The equilibria considered in the speciation calculations, the corresponding stability constants for the calcium-phosphate system, and the values of the thermodynamic solubility products at 25, 50, and 70 °C8,26 are summarized in Table 1. The experimental conditions were selected so that the working solutions were supersaturated with respect to HAP as may be seen from Figure 1A-C, in which the vertical bars indicate the experimental conditions in comparison with the solubility isotherms of the various calcium phosphate phases. An additional feature shown in Figure 1 is that some of the solutions were supersaturated with respect to all calcium phosphate phases. According to Ostwald’s rule, less stable phases as DCPD may form and even stabilize kinetically for shorter or longer time periods. With the progress of precipitation, however, the solution supersaturation with respect to DCPD is reduced, and the precipitation rate is decreased accordingly. The initially formed DCPD may function as a suitable substrate for the formation of OCP and HAP at supersaturation ratios significantly lower than those corresponding to spontaneous formation.16 The spontaneous formation of OCP and HAP is not possible after formation of DCPD since the calcium and phosphate ions consumption for the formation of DCPD causes a significant decrease in the respective supersaturation ratios. An alternative route for the formation of OCP and HAP is DCPD hydrolysis depending on the solution conditions.27 There are therefore two possible routes for the formation of most stable calcium phosphate phases: overgrowth and hydrolytic phase transformation. The design of the experiments in the present work aimed at the identification of the phase transformation sequence resulting in the formation of the thermodynamically most stable HAP.28 The calcium phosphate phases present in the solid precipitates formed were identified by powder XRD analysis using the appropriate reference materials.29 It should be noted that in the case of calcium phosphates, especially when precursor phases form and transform to more stable phases, the characterization of mixtures of phases is often difficult. Thus, additional phase characterization was done by SEM through the characteristic morphology of the mineral phases. The identification of OCPHAP mixtures and the determination of their respective contents (for their characterization as a minor or major phase) were based on the comparison of the reflections of the XRD patterns.30 The 211, 112 reflections of hydroxyapatite and the reflections 710, 420 of OCP are in close proximity. When however both phases are present in significant proportions, the diffraction pattern is distinct from that of either pure HAP or pure OCP.31,32 Typical examples of the XRD patterns obtained from OCP-HAP

mixtures are shown in Figure 2. As may be seen, the prevalence of one phase over the other clearly differentiates the pattern, allowing for safe conclusions concerning the mineral composition. It should be noted that the presence of OCP in the precipitated solids was also verified by FTIR and Raman spectroscopy.33-35 The differentiation between OCP and HAP may be done by examination of the bands in the range 500700 cm-1. Stabilization of transient phases including DCPD and OCP have been reported for the formation of calcium phosphate salts in soils in the presence of humic substances.36 The initial experimental conditions and calcium phosphate phases formed over the range of temperature, pH, and ionic strength values investigated are summarized in Tables 2-4. (i) Effect of Solution pH. At 25 °C, I ) 0.1 M, and in the absence of sand, DCPD was found to be kinetically stabilized in the pH range 6.0-7.0. A mixture of HAP, OCP (major phases) and DCPD (minor phase) was identified at pH ) 8.0. Formation of HAP is favored through fast hydrolysis of DCPD to OCP and HAP. HAP was not detected in the presence of sand at the same initial conditions, suggesting that the substrate acted as a stabilizing agent for the transient phases. A small amount of OCP was found at pH ) 8.0, but the main solid component was DCPD. The presence of OCP has been suggested to be due to the residual supersaturation, which is maintained at relatively high levels in our experimental conditions.37 At lower pH ( OCP and (b) OCP . HAP. Estimates are based on refs 29 and 37.

prevalent at the surface of the silicic sand grains favoring the formation of the acidic DCPD. At higher ionic strength values (0.5 M) and in the absence of sand, DCPD was found to precipitate in the solution at the pH range 6.0-7.0. Minor quantities of OCP were identified. At pH ) 8.0, however, the only calcium phosphate phase identified was HAP. At the same ionic strength, and in the presence of sand, DCPD was found to form almost exclusively in the solids precipitated at the pH range 6.0-7.0. At pH ) 8.0, HAP was found to be the main component of the solid precipitate, while small amounts of DCPD were also identified. It was attempted to identify the solid phase forming during the initial pH drop. Thus, past the initial pH drop the experiment was interrupted, and the solid was separated by filtration. The examination of the precipitate with XRD and SEM showed the formation of amorphous calcium phosphate (ACP), which rapidly converted to more stable phases depending on the solution conditions. The formation of ACP at similar conditions and its subsequent transformation to more stable calcium phosphate phases have been reported in the literature.39-42 The stability domain of the different calcium phosphate phases at room temperature and pressure conditions depends strongly on the pH of the aqueous environment. This is due to increasing supersaturation with respect to the basic calcium phosphates with increasing pH (Figure 1).37 The acidic DCPD is stabilized at pH values below pH 7, while the basic HAP is stable at higher pH values. In the intermediate pH range OCP is the dominant phase. On the basis of thermodynamics only, this is not anticipated, but due to kinetic factors. Less stable phases such as DCPD may still form, according to the reaction:

Ca2+ + HPO42- + 2H2O ) CaHPO4‚2H2O (s) Figure 1. Logarithm of the product of total calcium, TCa, and total phosphate, TP, concentrations, at equilibrium, with respect to DCPD, OCP, and HAP, at 25 °C (A), 50 °C (B), and 70 °C (C), as a function of pH. The bars correspond to the logarithm of the product of total calcium and total phosphate concentrations of the experiments of the present work.

Where dashed lines represent the surface of the solid substrate. As may be seen from the above equilibria, the silicate substrates upon immersion into the aqueous solution undergo a surface titration. Alternatively, it may be considered that low pH is

(8)

Reaction 8 is consistent with the initial changes in Ca2+ and H+ concentrations shown in Figure 4. As DCPD precipitates, the Ca2+ concentration decreases. Removal of HPO42- from the equimolar H2PO4- and HPO42- solution will make the solution more acidic due to reaction 9:

H2PO4- ) HPO42- + H+

(9)

Since thermodynamically DCPD is the least stable calcium phosphate salts at these conditions, other more stable phases



Table 2. Initial Experimental Conditions and Phases Formed at 25 °Ca pH

ionic strength M

sand grains substrate

6.0 6.0 6.0 6.0 7.0 7.0 7.0 7.0 8.0 8.0 8.0 8.0

0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5

no yes no yes no yes no yes no yes no yes

a

2.0

total calcium, Cat ) total phosphate, Pt mM 10.0 15.0

5.0

n.s. n.s. n.s. n.s. n.s. n.s. n.s. n.s. n.s. n.s. n.s.

n.s. n.s. n.s. n.s. DCPD, DCPA n.s. HAP, OCP

DCPD n.s.

DCPD DCPD

n.s. DCPD, DCPA DCPD, DCPA DCPD

n.s. DCPD, DCPA DCPD, DCPA DCPD, DCPA DCPD, DCPA OCP, DCPD, HAP DCPD, OCP HAP HAP, DCPD, DCPA

OCP, DCPD, HAP DCPD, OCP HAP HAP, DCPD

HAP

20.0 DCPD DCPD DCPD DCPD, DCPA DCPD, DCPA DCPD, DCPA DCPD, OCP DCPD, DCPA HAP, OCP, DCPD DCPD, OCP HAP, DCPD

Bold: major phase, n.s.: no solid formation. Table 3. Initial Experimental Conditions and Phases Formed at 50 °Ca

pH

ionic strength M


6.0 6.0 6.0 6.0 7.0 7.0 7.0 7.0 8.0 8.0 8.0 8.0

0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5


a


2.0

5.0

n.s. n.s. n.s. n.s.

n.s. n.s. n.s. n.s. HAP, OCP DCPD, OCP HAP OCP, HAP HAP HAP, OCP HAP OCP, HAP

n.s. n.s. n.s.

DCPD DCPD, DCPA n.s. n.s. HAP, OCP OCP, DCPA HAP HAP, OCP, DCPD HAP, OCP HAP, OCP HAP HAP, OCP

DCPD, OCP DCPD, DCPA DCPD DCPD HAP, OCP OCP, HAP, DCPD HAP HAP HAP HAP, OCP HAP HAP

20.0 DCPD, OCP DCPD, DCPA DCPD, OCP DCPD HAP, OCP OCP, HAP HAP HAP, OCP HAP HAP, OCP HAP HAP

Bold: major phase, n.s.: no solid formation. Table 4. Initial Experimental Conditions and Phases Formed at 70 °Ca

a

pH

ionic strength M


6.0 6.0 6.0 6.0 7.0 7.0 7.0 7.0 8.0 8.0 8.0 8.0

0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5 0.1 0.1 0.5 0.5


2.0

5.0

n.s. n.s. n.s. n.s.


n.s. n.s. HAP HAP, OCP HAP, OCP

HAP n.s. n.s.

HAP, OCP HAP, OCP HAP, OCP HAP

20.0

HAP, OCP

HAP

HAP, OCP

OCP, HAP

OCP, HAP OCP, HAP HAP, OCP OCP, HAP HAP, OCP OCP, HAP HAP, OCP HAP, OCP HAP, OCP HAP

OCP, HAP OCP, HAP HAP, OCP OCP, HAP HAP OCP, HAP

HAP, OCP HAP, OCP HAP, OCP OCP, HAP HAP, OCP HAP, OCP HAP HAP

HAP HAP, OCP HAP

Bold: major phase, n.s.: no solid formation.

may precipitate. Thus, HAP and OCP can form according to the reactions:

5Ca2+ + 3PO43- + OH- ) Ca5(PO4)3OH (s)

(10)

8Ca2+ + 6PO43- + 2H+ + 5H2O ) Ca8H2(PO4)6‚5H2O (s) (11) Both HAP and OCP give similar profiles in the Ca2+ and H+ concentration as already discussed for DCPD. About 40-50 min from the initiation of the precipitation process, the H+ concentration goes through an inflection point. Similar variations are observed for the Ca2+ concentration (first a minimum followed by a maximum). The latter observation of increase of the calcium ion activity is surprising for a precipitation process leading to ion consumption rather than release. Equally surprising is the behavior of the H+ concentration in the solution as a continuous increase is anticipated. A possible explanation for

the profiles shown in Figure 4 may be that dissolution and precipitation take place simultaneously.39 DCPD, the kinetically favored phase43 was precipitated and identified after 10 min from the onset of precipitation (see Figures 5 and 6a). It may be suggested that DCPD starts to dissolve when the supersaturation decreases due to the reduction in calcium phosphate and H+ concentrations according to reaction 8. As a result, a thermodynamically more favored phase, e.g., OCP, is likely to nucleate and grow further. The formation and persistence of DCPD in the precipitates was confirmed by XRD analysis as may be seen in Figure 5. Small OCP amounts were identified along with the principal DCPD component, as shown in Figure 6b. The presence of minor amounts of OCP, as may be seen in Figure 6a,b may be attributed to the stabilizing effect of the presence of Mg2+ ions, which may inhibit OCP hydrolysis.44-46 Dissolution of OCP and the subsequent formation of the stable HAP phase may be retarded and yield the concentration profile discussed above. It is interesting to note that in the precipitates


Figure 3. Variation of pH with time for the spontaneous precipitation of calcium phosphate phases at 25 (a), 50 (b), and 70 °C (c) in the presence of sand grains. Total calcium ) 15 mM, total phosphate ) 15 mM, initial solution pH ) 7.0, and ionic strength ) 0.1 M.

Figure 4. Variation of Ca2+ ion-selective electrode potential and of the solution pH from a combination glass/Ag|AgCl electrode versus time. The left-hand side ordinate shows the potential change of the H+ and the Ca2+ ion-selective electrodes measured past the mixing of the calcium (5 mM) and phosphate (10 mM equimolar HPO42- and H2PO4-) solutions in the absence of sand. The right-hand side ordinate shows pH measured from the calibrated pH electrode. The measurements were performed at about 20 °C. No NaCl was added for ionic strength adjustments.

formed at 25 °C the high-temperature phase DCPA was found by the XRD analysis even. Again this is a case of kinetic stabilization and not a drying effect of the specimens since this treatment was done at very mild conditions (air-drying at temperatures that did not exceed 60 °C). From the results presented, it may suggested that at pH values as high as 7.0, acidic calcium phosphates are favored. At higher pH values, however, the rapid phase transformation resulted into the gradual prevalence of phases relatively more stable than DCPD like OCP and HAP. The important role on the stabilization of transient acid phases was played by silicic sand grains. (ii) Effect of Temperature. Temperature is expected to affect both the equilibrium constants but most important the kinetics of precipitation. At 50 °C the stabilization of DCPD was found to occur at lower pH. DCPD was thus formed only in supersaturated solutions with pH ) 6.0 in the absence of sand, while small amounts of OCP were also identified in the precipitate. The formation of DCPD was favored at solution pH ) 7.0 only in the presence of sand grains. This phase however rapidly hydrolyzed to OCP, and finally DCPD was detected only as a minor phase component in the solid

Hafez et al.

Figure 5. XRD of DCPD (1) and DCPA (2) precipitated from calcium phosphate supersaturated solutions. Solids were sampled after 10 min from the beginning of the experiment, the results of which are presented in Table 2.

precipitate. In the absence of sand grains and in supersaturated solutions at pH ) 7.0, the solids formed consisted of mixtures of HAP and OCP. In the precipitates formed at pH ) 8.0 and in the absence of sand grains, HAP was the only solid phase identified. In the presence of sand grains, the presence either of DCPD or OCP even in minor contents suggested the formation and transformation of DCPD. At solution pH ) 7.0, OCP was identified as major phase, while DCPD and HAP were detected in small quantities. It is possible that the DCPD crystals nucleating and growing on the sand grains hydrolyzed to the more stable HAP. It may be suggested that the initial formation of DCPD at these conditions may be favored by the low pH at the surface of the sand grains as already mentioned. It should be noted that the precipitation on the sand grains took place without any induction times. HAP and OCP were identified as the major and minor phases in the precipitates formed in solutions at pH ) 8.0, respectively. The formation of HAP at pH higher than 6.0 and the absence of DCPD and OCP may be attributed to the fast hydrolysis of these phases to HAP. The hydrolysis of DCPD to OCP and OCP to HAP is significantly favored with temperature.36,38 At the same initial conditions and in the presence of sand, HAP was formed at solution pH ) 8.0, while small quantities of OCP were also detected in the precipitates. OCP was identified as major phase and HAP as minor phase at solution pH ) 7.0 and at low initial calcium and phosphate concentrations. The higher the supersaturation at the same pH (provided that the solution pH is not acidic), the more rapid were the precipitation and conversion of the metastable phases to the thermodynamically stable HAP. The most distinguishable phase, although in small amounts with respect to the overall solid composition, was OCP, while in the acid pH region DCPD was the only phase identified in the solid precipitate. At 70 °C, DCPD was not detected in the solid precipitates in the pH range 6.0-8.0. In these experiments, HAP and OCP were the only crystalline phases present. At low ionic strength and in the absence of sand, HAP was identified as a major phase and OCP as a minor phase at all pH values tested. In the presence of sand, OCP was the main component of the mineral phase precipitated at pH ) 6.0. For pH > 6.0, however, HAP was the principal solid component, suggesting fast hydrolysis of the transient OCP. In the cases in which the precipitation



Figure 6. Scanning electron micrographs of precipitated material, which was sampled 10 min after mixing of the calcium chloride and phosphate solutions. The area A, in panel a, is shown on the right-hand side (panel b), with double magnification. A different phase, probably OCP, seems to be present in this area.

took place without the mediation of the sand at pH ) 6.0 and ionic strength of 0.5 M, OCP was found to be the main component. At higher pH values, a mixture of HAP and OCP was obtained, the former coming apparently from the hydrolysis of transient phases. The precipitates obtained at 70 °C consisted almost exclusively of HAP. This may be ascribed to the rapid hydrolysis of the transient phase OCP, which forms first. At the same temperature and at ionic strength 0.5 M corresponding to relatively lower supersaturation at pH ) 7.0, the metastable OCP was stabilized. At higher pH values, corresponding to drastically increased solution supersaturation with respect to HAP (see Figure 1), the only mineral phase identified in the solid precipitate was HAP. It is interesting to note that DCPA was identified in the solids formed up to 50 °C. At 70 °C, no DCPA was identified despite the fact that the driving force was favorable. Apparently at higher temperatures and hence higher supersaturations the rates of transformation are significantly faster, eliminating any trace of transient phases forming. The present work shows that the calcium phosphate phases formed from supersaturated solutions either in the presence or in the absence of foreign substrates depend on the solution supersaturation. The formation of OCP is significantly favored at relatively higher temperatures (higher than 50 °C) in agreement with literature data.15 Moreover, it should be noted that OCP was stabilized at lower pH values as the temperature increased.28 This stabilization may be attributed to solubility changes concomitant with the temperature increase. The phase formation and transformation of the calcium phosphate system may be visualized by the pH-time profiles shown in Figure 3. As may be seen, at 25 °C the solution pH profile, past an initial fast drop, showed a plateau region followed by a second stage of pH drop. This inflection in the pH-time profile may be attributed to transformations of transient calcium phosphate phases to the most stable HAP.29 Alternatively, the decrease in pH may be attributed to the overgrowth of calcium phosphates on the phases nucleated initially through kinetic stabilization. A similar pH-time profile, though less pronounced, was obtained in the experiments done at 50 °C. Upon further increase in temperature to 70 °C as may be seen in Figure 3, the inflections were markedly reduced in comparison with 25 and 50 °C, suggesting rapid phase transformations. At 25 °C, the thermodynamically less stable polymorphs are kinetically stabilized at the initial stages of precipitation. As the temperature

increased the kinetic stabilizing effects are reduced and the formation of thermodynamically more stable phases are favored. The OCP phase was identified from XRD and from morphological analysis (Figure 6b). It might be suggested that at 20-50 °C, the rate of crystallization of DCPD is faster than the other phases at high supersaturations. DCPD however is not the thermodynamically most stable phase. It may therefore dissolve when the solution supersaturation drops. OCP is a more stable phase, but the corresponding rate of crystallization is slower. It may therefore take some time before this phase starts to form. This may explain why mainly DCPD is present with minor quantities of OCP. The pH-time profiles shown in Figure 3 showed the same trend as those observed in Figure 4. The minimum in the pH value, however, occurs earlier at higher temperatures in agreement with the fact that higher nucleation and precipitation rates are expected for all calcium phosphate phases at higher temperatures. It was therefore reasonable to suggest that the precipitation mechanism is the same for all the experimental conditions investigated. The relative stability of the mineral phases forming plays an important role in the functionality of the oil well formations not only from the point of view of the widely different solubility of the various calcium phosphate phases but also from their different crystallographic and morphological characteristics. Below micron to nanosized HAP resulting from the transformation of transient phases may block the pores of the oil formation reducing permeability of the sand packs. Preliminary measurements of this parameter in our laboratory corroborate this suggestion. The platy formations on the other hand of OCP (1-3 µm) appear to ensure optimal permeability combined with formation stabilization.6,7 The formation of large DCPD platy crystals (20-30 µm) seemed to yield more effective consolidation, but the formation permeability to fluids is not satisfactory. (iii) Effect of Ionic Strength. The experiments in the present work were conducted at relatively low values of ionic strength. We have avoided working at higher values of ionic strength perhaps closer to the chemical environment in the oil wells because at this stage of research we would like to obtain a better understanding of mechanistic effects. We have thus preferred to use solutions with ionic strength values, which could be handled with the same principles and with the same equilibrium computations. From the results summarized in Tables 2-4, it may be suggested that the ionic strength of the supersaturated


solutions affected precipitation to the extent that it influenced the solution supersaturation. Lower supersaturations result in faster transformation of any transient phases that may be formed and short lived.47 At low temperatures, kinetics factors favored the formation and stabilization of the thermodynamically less stable phases. With increasing temperature the most stable phases were formed. The nature of the solid phases precipitated and identified in the present work are in good agreement with our earlier report according to which at all supersaturations the least stable calcium phosphate phase is precipitated first, transforming through the formation of the intermediate OCP to the thermodynamically most stable HAP.7 The stabilization of the transient phases however is kinetically controlled. The results of the present work suggest that it is possible to tune the chemical environment in which consolidation takes place, in such a way that a given phase may have a preference for precipitation. Fine-tuning includes in the first place adjustment of local supersaturation. This may be achieved through components concentrations, solution pH, ionic strength, and temperature adjustment. The relevance of the present work to the oil field situation may be drawn from the dependence of the various calcium phosphate phases forming on the silica particles on the solution temperature. Supersaturation is strongly affected by both temperature and ionic strength. These parameters are often quite variable depending on the characteristics of the individual oil reservoirs. However, the results of the present work may be applicable for the prediction of calcium phosphate scale formation and of the formation of various phases provided that the results are translated into supersaturations, which relate to thermodynamic parameters such as the Gibbs free energy for the formation of calcium phosphates from supersaturated solutions. Conclusions In supersaturated solutions and in the absence of extraneous material, the formation of DCPD was favored at low temperatures at pH values below 8.0. At pH ) 8.0, the formation of HAP was favored through hydrolysis of the thermodynamically less stable calcium phosphates. This transformation was confirmed from the pH variation with time. As temperature increased, the formation of the thermodynamically most stable phases, HAP and OCP, was favored at pH > 6.0, while the formation of DCPD was favored at lower pH values. At 70 °C, the thermodynamically more stable phase HAP was predominant. It is interesting that with the exception of the highest temperature tested (70 °C) small amounts of DCPA were identified in the precipitates. The stabilization of this unstable, high-temperature transient phase was favored by the presence of sand grains as in the case of DCPD. The pH limit for the formation of OCP decreased as temperature increased. Thus, at 25 °C and pH ) 8.0, OCP and HAP were identified in the solid precipitates. At 25 °C and in the absence of sand, the transformation of DCPD to OCP was accelerated. The transformation proceeded finally to the formation of the thermodynamically most stable HAP. In supersaturated solutions of calcium phosphate at higher temperatures and in the absence of sand, HAP was the solid phase formed almost exclusively. The presence of a few OCP crystals however in the precipitate suggested that the mineral phase identified as HAP could be the product of rapid transformation of transient mineral precursor phases. At 70 °C, the formation of OCP was identified as a minor component phase even at pH ) 6.0. At higher pH values (7.0 or 8.0), the rapid transformation of the transient phases

Hafez et al.

resulted in the formation of HAP as the principal component of the solid precipitate. The pH-time profiles during the course of precipitation of calcium phosphates showed local maxima and minima, suggesting hydrolysis of less stable phases and crystallization of more stable calcium phosphates. During the initial pH drop, amorphous calcium phosphate was formed, which rapidly converted to more stable phases. Seeding the supersaturated solutions with silicate sand at all temperatures investigated resulted in the formation of calcium phosphate crystals on the seed particles without the mediation of measurable induction times. At 25 °C, the precipitate consisted of DCPD crystals. The surface acidity of the silicate substrate may be responsible for the development of a favorable local environment for the deposition of the acid calcium phosphate phase. At 50 °C DCPD, though as a minor component, was stabilized even at pH ) 7.0, coexisting with OCP. At 70 °C and at pH 6.0 and 7.0, the dominant phase precipitated consisted mainly of OCP crystals identified together with the most stable HAP. Finally, it should be noted that the effect of ionic strength was limited to the respective effect on the solutions supersaturation. Acknowledgment. This work has been supported financially by the European Social Fund (ESF), Operational Program, for Educational and Vocational Training II (EPEAEK) and particularly the Program IRAKLEITOS (HERACLITUS), the Institute of Chemical Engineering and High Temperature Chemical Processes (ICE/HT-FORTH), the Norwegian University of Science and Technology and the Norwegian Research Council. The authors are grateful for the financial support. References (1) Penberthy, W. L.; Shaughnessy, C. M. Sand Control, 1st ed.; Society of Petroleum Engineers: Richardson, TX, 1992. (2) Sofitech NV, BE, European Patent 0 721 050 A2, 1996. (3) Mobil Oil Corp., European Patent 0 476 820 A2, 1992. (4) Shell International Research, European Patent 0 030 753 A2, 1982. (5) Mobil Oil Corp., European Patent 0 414 431 A2, 1991. (6) Paraskeva, C. A.; Charalampous, P. C.; Stokka, L. E.; Klepetsanis, P. G.; Koutsoukos, P. G.; Read, P.; Østvold, T.; Payatakes, A. C. J. Colloid Interface Sci. 2000, 232, 326-339. (7) Gustavsen, Ø.; Larsen, T.; Østvold, T.; Paraskeva, C. A.; Hafez, I. T.; Toliza, A.; Klepetsanis, P. G.; Koutsoukos, P. G.; Payatakes, A. C. Sand Stabilization with Mineral Precipitation, In Progress in Mining and Oilfield Chemistry, Vol. 5. AdVances in Incremental Petroleum Production; Lakatos I., Ed.; Akademiai Kiado: Budapest, 2003; Vol 5, pp 259-278. (8) Valsami-Jones, E. Mineral. Mag. 2001, 65, 611-620. (9) Meyer, J. L.; Eanes, E. D. Calcified Tissue Res. 1978, 25, 209-216. (10) LeGeros, R. Z.; Kijkowska, R.; LeGeros, J. P. Scanning Electron Microsc. 1984, 1771-1777. (11) Eanes, E. D.; Termine, J. D.; Nylen, M. U. Calcified Tissue Res. 1973, 12, 143-158. (12) Nancollas, G. H. In vitro studies of calcium phosphate crystallization, In Biomineralization; Mann, S., Webb, J., Williams, R. J. P., Eds.; VCH Publishers: Weinheim, 1989; pp 157-187. (13) Nancollas, G. H.; Amjad, Z.; Koutsoukos, P. G. Calcium Phosphates. Speciation, Solubility and Kinetic Considerations. In Chemical Modelling in Aqueous Systems, ACS Symposium Series, No. 93, Jenne E. A.; Ed.; American Chemical Society: Washington DC, 1979; pp 475-497. (14) Tsuge, H.; Yoshizawa, S.; Tsuzuki, M. Chem. Eng. Res. Des. 1996, 74, 797-802. (15) Fernadez, E.; Gil, F. J.; Ginebra, P.; Driessens, F. C. M.; Planell, J. A.; Best, S. M. J. Mater. Sci. 1999, 10, 177-183. (16) Freche, M.; Heughebaert, J. C. J. Cryst. Growth 1989, 94, 947954. (17) Kim, C. Y.; Clark, A. E.; Hench, L. L. J. Biomed. Mater. Res. 1992, 26, 1147-1161. (18) Filgueiras, M. R.; La Torre, G.; Hench, L. L. J. Biomed. Mater. Res. 1993, 27, 445-453. (19) Kibalczyc, W.; Sokolowski, T.; Wiktorowska, B. Cryst. Res. Technol. 1984, 19, 27-32.

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Calcium Phosphate Overgrowth on Silicate Sand - ACS Publications

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