Calcium Carbonate Crystallization in the Presence of Aspartic Acid P.
Malkaj‡
and E.
Dalas*,†
Department of Chemistry, University of Patras, GR-26504 Patras, Greece, and Department of Physics, University of Gjirokastra, Gjirokastra, Albania
CRYSTAL GROWTH & DESIGN 2004 VOL. 4, NO. 4 721-723
Received March 26, 2003
ABSTRACT: The kinetics of vaterite (CaCO3) crystallization on calcite in the presence of aspartic acid was investigated by the constant composition technique. The presence of aspartic acid in the supersaturated solution stabilizes this calcium carbonated polymorph possibly through the development of active growth sites, which should show chemical and structural affinity to this mineral phase. The number of ions forming the critical nucleus was n* ) 4, and a surface energy of 31 mJ m-2 for the growing phase was estimated. The apparent order was found to be 1.75 ( 0.25, indicative for a surface diffusion controlled mechanism. Introduction Calcium carbonate is the most abundant mineral formed by organisms, followed by silica, calcium phosphate, and others.1 Biomineralization results in complex materials that are characterized by a remarkable level of molecular control of the particle size, structure and morphology, aggregation, and crystallographic orientation of the calcium carbonate phases.2-4 The variety of biomineralizates can be expressed by the fact that approximately 128 000 species of molluscs exist. The majority of them (Conchifera) produce shells and capsules from calcium carbonate having different sizes, shapes, and color patterns.5 Many shells are constructed within frameworks that may be lamellar, columnar, or reticular. In each case, there are two important features of assembly. First, a relatively inert structural frame is built from insoluble macromolecules.6,7 Second, acidic proteins, such as aspartic acid, are assembled on the hydrophobic scaffold. Mineralization takes place at the interface between the acidic proteins and the aqueous environment.4,8,9 Also, the oyster shell matrix contains 20-30 mol % aspartate.1 The aim of the present work is to investigate the effect of aspartic acid on the growth of carbonate crystals by the constant composition technique10,11 and to attempt to answer the following question: Does aspartic acid affect the nature, the rate, or the particle size of the calcium carbonate forming phases? The methodology was applied for studying the crystallization process of calcium carbonate in seeded growth and spontaneous precipitation experiments because it permits accurate assessing of the rates of crystallization and the nature of the precipitating crystalline polymorphs.10,11 Experimental Section The experiments were performed in a thermostated doublewall Pyrex vessel at 25.0 ( 0.1 °C. The total volume of the working solution was 0.200 dm3, and the stable supersaturated * Corresponding author. Tel: +30-2610-997145; fax: +30-2610997118. E-mail:
[email protected]. † University of Patras. ‡ University of Gjirokastra.
solutions employed were prepared by mixing equal volumes of calcium nitrate and sodium bicarbonate solutions, made from standardized stock solutions as described in detail elsewhere.12 The pH of the solutions was measured by a glass/ saturated calomel electrode (Metrohm), standardized before and after each experiment with NBS buffer solutions.13 The pH of the working solution was adjusted to 8.5 by the addition of standard potassium hydroxide (Merck, Titrisol) and the pH remained constant for at least 10 h, indicating the lack of any precipitation in the solution. Prior to the pH adjustment, a quantity of 2-40 mg of aspartic acid was added to the working solution. Aspartic acid was purchased from Aldrich and used without further purification. The calcium carbonate precipitation process was started after the lapse of well-defined induction periods for the spontaneous precipitation experiments or immediately after the introduction of 100 mg of synthetic calcite crystals, prepared according to the method described by Reedy and Nancollas14 and characterized by powder X-ray diffraction (Phillips PW 1830/1840 using CuKR filter), FT-IR spectroscopy (Perkin-Elmer 16-PC FT-IR using KBr pellets), and thermogravimetric analysis (Du Pont 910). The specific surface area of the seed crystals, as determined by a multiple-point BET method, was found to be 3.2 m2 g-1. A pH drop as small as 0.005 pH units was sufficient to trigger the addition of titrant from two mechanical burets of an appropriately modified pH stat (Metrohm Dosigraph with 614 Impulsomat). The two burets contained calcium nitrate and sodium carbonate titrants having the stoichiometry of the precipitating salt (calcium/carbonate ) 1:1).7,11,12 By inclusion in the titrants of the appropriate calcium nitrate concentration, as well as the appropriate potassium nitrate concentration for the ionic strength, the initial conditions could be maintained constant throughout the precipitation process. Calcium nitrate, potassium nitrate, sodium bicarbonate, and sodium carbonate were purchased from Merck (pro analysis). The crystal growth rates, R, were easily and accurately obtained from the recorder traces of titrants (corresponding to the moles of calcite precipitating) as a function of time.10,11 The constancy of the solution composition was checked by regular sampling, filtering the samples through membrane filters (0.22 µm Gelman) and analyzing the filtrates for calcium by atomic absorption spectroscopy (Varian 1200). At the end of the experiment, the solids were collected by filtration and examined by scanning electron microscopy (SEM), FT-IR spectroscopy, powder X-ray diffraction, and thermogravimetric analysis (TGA).
10.1021/cg030014r CCC: $27.50 © 2004 American Chemical Society Published on Web 06/15/2004
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Malkaj and Dalas
Results and Discussion The driving force for the CaCO3 formation is the change of Gibbs free energy, ∆G, occurring during the transfer from the supersaturated solution to equilibrium. Thermodynamically speaking, the criterion for a physical process to take place is the change in the Gibbs free energy between the initial and the final states. If the change of the Gibbs free energy is found to be negative, then the process is spontaneous (i.e., ∆G < 0). It is known that the difference in Gibbs free energy refers to the corresponding change of chemical potentials:
∆Gx ) ∆µ ) µ(sat) - µ(supersat)
(1)
Table 1. Crystal Growth of Vaterite on Calcite in the Presence of Aspartic Acid at Sustained Supersaturation, 25 °C, pH 8.50a Cat (10-3 mol dm-3)
aspartic acid (mg)
ionic strength (10-2 mol dm-3)
∆Gvat (kJ mol-1)
R (10-3 mol min-1 m-2)
3 2.75 2.5 2.25 2 3 3 3
20 20 20 20 20 0 2 40
7.19 6.60 5.00 4.5 4.0 7.19 7.19 7.19
-1.46 -1.30 -1.10 -0.89 -0.65 -1.46 -1.46 -1.46
8.14 6.66 3.66 2.99 1.82 5.71 7.00 10.09
a Total calcium (Ca ) ) total carbonate (C ), 0.5 mg of calcite t t powder/mL.
For the crystal growth of CaCO3 the chemical potential in the two states are
µ(sat) ) µ0 + RgT ln[aCa2+(sat)aCO32-(sat)]1/2
(2)
µ(supersat) ) µ0 + RgT ln[aCa2+(supersat)aCO32-(supersat)]1/2 (3) where a(ion) is the ionic activity, Rg is the gas constant, and T is the absolute temperature. The difference between the chemical potential in the two states is
[
∆µx ) RgT ln
]
1/2
aCa2+(sat)aCO32- (sat)
aCa2+(supersat)aCO32- (supersat)
)
[ ]
RgT ln
K 0sp,x (4) IPx
0 In eq 4, IPx is the activity product, and K sp,x is the thermodynamic solubility product of the polymorph x. After substitution in eq 1, the Gibbs free energy change for the crystallization of CaCO3 is given by
∆Gx ) -RgT ln
( ) IPx
0 K sp,x
1/2
)-
RgT ln Ωx 2
(5)
where x denotes the precipitating polymorph x and Ωx is the supersaturation ratio. Spectroscopic examination of the precipitating solids in all cases by (a) X-ray diffraction15,16 exhibit the characteristic reflections for vaterite hkl: 110, 111, 112, 300, 302, 114, 222; (b) FTIR spectroscopy (exhibit the characteristic absorptions for vaterite at 1480, 1070, 873, and 745 cm-1);17-19 and (c) SEM photographs typical for vaterite spherical shape crystals. The absence of hydrated polymorphs was also ruled out by differential scanning calorimetry.20 The solution speciation in all experiments was computed from pH, the total calcium balance, and the electroneutrality conditions, assuming a system in which the partial pressure of CO2 was kept constant. This assumption is valid11 because the solution’s pH is high and the volume of the air over the working solution 0 ) 1.222 × 10-8 (the is minimal. From eq 5 and K sp,v thermodynamic solubility of vaterite21), the driving force ∆Gv for the formation of vaterite was calculated. The initial conditions of the experiments reported herein are summarized in Table 1. The measured rates were
Figure 1. Plot of the logarithm of the initial rates, R, as a function of the logarithm of the initial calcium concentrations.
typical for crystal growth at sustained supersaturation,10,11 and their reproducibility was better than 2% (a mean of five experiments). From the slope of the plot of the logarithm of the initial concetration,22 the critical size of the nucleus, n*, may be estimated:
d ln R/d ln[Ca2+]0 ) n*
(6)
Such a plot is shown in Figure 1 and a value of n* ) 4 was obtained. The same value was obtained for the vaterite formation on cholesterol17 as well as for the overgrowth of vaterite on calcite seed crystals in the presence of glutamic acid.23 Using nucleation rate equations derived from the classical homogeneous nucleation theory, interfacial energy σ, for the vaterite overgrowth was calculated:
[
R ) K exp -
Φb Vm2 σ3 k3T3(ln Ωv)2
]
(7)
in which Vm is the molecular volume of the precipitated vaterite (3.129 × 10-29 m3), k is the Boltzmann constant, K is another constant, and b is the shape factor (b ) 16.76, assuming spherical shape nuclei and Φ a constant (0 < Φ < 1).6,7,19,23 A plot of ln R against (1/ln Ωv)2 according to eq 7 results in a straight line from the slope of which a value of 31 mJ m-2 was obtained for the surface energy of growing vaterite. Similar value for vaterite formation has been published and summarized
Calcium Carbonate Crystallization
Crystal Growth & Design, Vol. 4, No. 4, 2004 723
ber of ions forming the critical nucleus was found to be n* ) 4 and a surface energy of 31 mJ m-2 was estimated from kinetics data, a relatively low value that may be attributed to the heterogeneous character of vaterite nucleation. The consideration of the polarity of the CdO bond in which the negative charge is shifted toward the oxygen atom30 suggests that the positively charged calcium ions were attracted by the electric field. The electrostatic interactions temporarily fixed Ca2+ ions near CdO groups that along with the diffused CO22- ions toward the fixed Ca2+ may initiate the critical nuclei for vaterite formation. The apparent growth order was found to be 1.75 ( 0.25 typical for a surface diffusion controlled growth mechanism. References Figure 2. Rate of crystallization of vaterite on calcite seed crystals in the presence of 20 mg of aspartic acid. Table 2. Surface Energy σ for Vaterite Formation conditions
σ (mJ m-2)
on cholesterol on carboxylated copolymer on fibrin on calcite in the presence of glutamic acid on lysine spontaneous precipitation by the free drift method spontaneous precipitation in the presence of chondroitin sulfate on calcite in the presence of aspartic acid
11 24 21 12
17 18 19 23
33 34
25 24
52
26
31
this work
ref
in Table 2. The theoretical value for the surface energy computed for homogeneous nucleation24 is σ ) 90 mJ m-2. The high value predicted above is pertinent to homogeneous nucleation conditions as contrasted to our experiments, where the new phase is grown on a foreign substrate. The rate of crystallization, R, is22
R ) k′ S nv
(8)
In eq 8, k′ is the apparent specific rate constant for the crystallization process, n is an apparent growth order, and Sv is the relative solution supersaturation (Sv ) Ωv1/2 - 1). A plot of ln R ) f(ln Sv) according to eq 8 yielded a straight line, from the slope of which an apparent order of n ) 1.75 ( 0.25 was calculated as shown in Figure 2 pointing to a surface diffusion controlled spiral growth mechanism. Similar results were obtained for the crystallization of vaterite on cholesterol,17 on carboxylated copolymer,18 and on calcite in the presence of glutamic acid.26 Different amounts of aspartic acid added to the supersaturated solutions resulted in an acceleration effect of the crystal growth process. Analogous results were obtained in the presence of glutamic acid26 and of benzotriazoles.27 In conclusion, it can be said that the presence of aspartic acid in the supersaturated solutions stabilizes the vaterite polymorph possibly through the development of active growth sites, which show chemical and structural affinity to this mineral phase.28,29 The num-
(1) Sikes, C. S.; Wheeler, A. P. CHEMTECH 1988, October, 620. (2) Mann, S.; Archibald, D. D.; Didymus, J. M.; Douglas, T.; Heywood, B. R.; Meldrum, F. C.; Reeves, N. J. Science 1993, 261, 1286. (3) Mann, S. New Sci. 1990, 10, 42. (4) Mann, S. Nature 1993, 365, 499. (5) Krampitz, G.; Graser, G. Angew. Chem., Int. Ed. Engl. 1988, 27, 1145. (6) Manoli, F.; Koutsopoulos, S.; Dalas, E. J. Cryst. Growth 1997, 182, 116. (7) Manoli, F.; Dalas, E. J. Cryst. Growth 1999, 204, 369. (8) Borman de Jong, A. H.; Huizinga, M.; Kok, D. J.; Westbroek, P.; Bosh, L. Eur. J. Biochem. 1982, 129, 179. (9) Addadi, L.; Berman, A.; Oldak, J. M.; Weiner, S. Connect. Tissue Res. 1989, 21, 127. (10) Tomson, M. B.; Nancollas, G. H. Science 1977, 200, 1059. (11) Kazmierczak, T. F.; Tomson, M. B.; Nancollas, G. H. J. Phys. Chem. 1982, 86, 103. (12) Giannimaras, E.; Koutsoukos, P. G. J. Colloid Int. Sci. 1987, 116, 423. (13) Bates, R. G. Determination of pH. Theory and Practice; Wiley: New York, 1973. (14) Reddy, M. M.; Nancollas, G. H. J. Colloid Int. Sci. 1971, 36, 166. (15) American Society for Testing and Materials, card file no. 25-127. (16) JCPDS-International Centre for Diffraction Data, card file no. 13-0192. (17) Dalas, E.; Koutsoukos, P. G. J. Colloid Int. Sci. 1989, 127, 273. (18) Dalas, E.; Klepetsakis, P. G.; Koutsoukos, P. G. J. Colloid Int. Sci. 2000, 224, 56. (19) Kanakis, J.; Dalas, E. J. Cryst. Growth 2000, 219, 277. (20) Dalas, E.; Kallitsis, J.; Koutsoukos, P. G. J. Cryst. Growth 1988, 89, 287. (21) Plummer, N. L.; Wigley, T. M. L.; Parkhurst, D. C. Am. J. Sci. 1978, 278, 179. (22) Nyvlt, J.; Sohnel, O.; Matuchova, M.; Broul, M. The Kinetics of Industrial Crystallization; Elsevier: Amsterdam, 1985; pp 301, 284-286. (23) Manoli, F.; Dalas, E. J. Cryst. Growth 2001, 222, 293. (24) Kralz, D.; Brecevic, L.; Nielsen, A. E. J. Cryst. Growth 1990, 104, 793. (25) Manoli, F.; Kanakis, J.; Malkaj, P.; Dalas, E. J. Cryst. Growth 2002, 236, 363. (26) Manoli, F.; Dalas, E. J. Cryst. Growth 2000, 217, 416. (27) Zafiropoulou, A.; Dalas, E. J. Cryst. Growth 2000, 219, 477. (28) Didymus, J. M.; Mann, S.; Benton, W. J.; Collins, I. R. Langmuir 1995, 11, 3130. (29) Sims, S. D.; Didymus, J. M.; Mann, S. J. Chem. Commun. 1995, 10, 1031. (30) Malkaj, P.; Chrissanthopoulos, A.; Dalas, E. J. Cryst. Growth 2002, 243, 233.
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