Langmuir 1985, 1 , 119-122 tradiol in 50% v/v methanol/water 0.1 M buffer. However, when the percentage of methanol was reduced to 5% v/v, an adsorptive couple was observed even a t “pH” 4.0. Cyclic voltammograms of a-estradiol (Figure 16B), the stereoisomer of &estradiol in the 17-OH position, under identical conditions to the P-estradiol experiments, showed significant differences in position and shape of the adsorptive responses relative to that of @-estradiol(Figure 16A). Oestrone (a phenolic steroid not possessing a 17-OH group) did not show any peaks at all in CV. Obviously the addition of the 17-OH group played an important role in the manifestation of the capacitative phenomena observed in CV as did the 3-OH group of the “A” ring phenol.
Conclusions The electron-transfer reactions of acetylenic steroid hormones at the mercury electrode are strongly influenced by interaction with the electrode surface. The surface interactions in turn are dependent upon the type and orientation of the substituents. The dichotomy in behavior between the oxidation process for the “phenolic” acetylenic steroid ethynylestradiol and the non-phenolic acetylene steroids can be rationalized in terms of their adsorptive properties a t the mercury electrode. Phenolic steroids in general have specific reversible interactions with the mercury surface a t high “pH”. These interactions are presumably a result of specific orientations on the electrode surface and depend upon the surface charge, stereoisomerism of 17-OH group,
119
phenol/phenolate A ring/”pH”, and methanol content. Such physical and chemical considerations have been expressed for specific orientations of other aromatic compounds a t the platinum e l e ~ t r o d e . ~ lFurthermore -~~ it is likely that these specific orientations for phenolic steroids facilitate the electrode process for the oxidation of mercury in the presence of ethynylestradiol, resulting in a comparatively well-defined process for the latter. However, at pH >lo, the non-phenolic acetylenic steroids do not exhibit specific electrode adsorption behavior and are much more strongly adsorbed than the corresponding phenolic steroids. The preferred orientation and strong adsorption of this class of steroid is believed to induce the observed irreversible electrode reaction behavior. The orientation of adsorbed intermediates is known to exert profound effects on electrochemical pro~esses.~~ Registry No. Ethynylestradiol, 57-63-6;mestranol, 72-33-3; norgestrel, 6533-00-2;chloronorgestrel, 14115-33-4;p-estradiol, 50-28-2;a-estradiol,57-91-0;testosterone,58-22-0; lynestranol, 52-76-6. (21) Soriaga, M. P.; Hubbard, A. T. J . Am. Chem. SOC.1982, 104, 2735-2742. (22) Soriaga, M. P.; Hubbard A. T. J. Am. Chem. SOC.1982, 104, 2742-2147. (23) Soriaga, M. P.; Hubbard, A. T. J. Am. Chem. SOC.1982, 104, 3937-3943. (24) Soriaga, M. P.; Stickney,J. L.; Hubbard, A. T. J. Mol. Catal. 1983, 21,211-221.
Crystal Growth of Calcium Phosphates in the Presence of Magnesium Ions M. H. Salimi, J. C. Heughebaert,*t and G . H. Nancollas Chemistry Department, State Uniuersity of New York at Buffalo, Buffalo, New York 14214 Receiued August 28, 1984 The influence of magnesium ions upon the crystallization rates of dicalcium phosphate dihydrate and octacalcium phosphate has been investigated at constant supersaturation. While having no detectable effect on the growth of dicalcium phosphate dihydrate, magnesium ions appreciably retard the rate of octacalcium phosphate growth, probably by adsorption at active growth sites. The mineralization of the thermodynamically most stable hydroxyapatite is much more strongly inhibited by the presence of these added ions, which may therefore mediate in the precipitation process by selectively stabilizing the more acidic precursor phases.
Introduction The precipitation of calcium phosphates is important both in biological mineralization and in natural water systems.’S2 Traces of magnesium ion have been shown to reduce the overall rate of seeded calcium phosphate crystallization3 and markedly delay the transformation of amorphous calcium phosphates to more stable apatitic phase^.^ In natural water systems, magnesium ion concentrations may be as high as 5 X mol L-1,2while in biological calcification magnesium concentrations ranging from 0.5% in outer tooth enamel layers to 2% in the innermost dentines is likely to have important consequences on the rates of remineralization. Numerous studies have On leave from Institut National Polytechnique de Toulouse, 31400 Toulouse, France.
been made of the influence of magnesium ions on calcium phosphate formati~n.~?’ The results of spontaneous precipitation experiments have suggested that magnesium ions kinetically hinder the nucleation and subsequent growth of hydroxyapatite (Ca5(P04)30H,hereafter HAP) by competing for lattice sites with the chemically similar but (1)‘Fological Mineralization and Demineralization”;Nancollas, G. H., Ed., Springer-Verlag: Heidelberg, 1982. (2) Stu“, W.; Morgan, J. J. “AquaticChemistry”;Wiley Interscience: New York, 1981. (3) Nancollas, G. H.; Tomazic, B.; Tomson, M. B. Croat. Chim. Acta 1976, 48,431. (4) Eanes, E. D.; Ratner, S. L. J. Dental Res. 1981, 60, 1719. (5) Shaw, J. H.; Yen, P. K. J. J. Dental Res. 1972, 51, 95. (6) Martens, C. S.: Harriss, R. C. Geochim. Cosmochim. Acta 1970,34, 621. (7) Ferguson, J.; McCarty, P. L. Enuiron. Sci. Technol. 1971,5, 534.
0743-7463/85/2401-0119$01.50/0 0 1985 American Chemical Society
120 Langmuir, Vol. I , No. 1, 1985
larger calcium ions.7 It was shown that in the presence of magnesium ions, magnesium-containing 6-tricalcium phosphate (Whitlockite, TCP) was f ~ r m e d . ~ , ~ In most of the studies, the calcium phosphate solutions were usually supersaturated with respect to all calcium phosphate phaseslOJ1in order of increasing solubility: HAP, TCP, octacalcium phosphate [Ca4H(P04)3.2.5Hz0; hereafter OCP] , and dicalcium phosphate dihydrate (CaHP04.2H20,hereafter DCPD). The development of the seeded growth techniques3J2J3enabled the rates of crystallization to be studied under highly reproducible conditions, and the use of a pH-stat to control the concentration of hydrogen ion during cry~tallization'~ limited complications arising from transformation of calcium phosphate phases during the reactions. However, in these experiments, the calcium and phosphate ionic concentrations decreased appreciably during the reactions, enabling the formation of different phases. Some of these problems were overcome by the development of a constant-composition technique in which the supersaturation was held constant throughout the precipitation process.15J6 The method has been successfully used to investigate the crystal growth of DCPD,17 OCP,lE,and HAP.16J7 In the present work, the constant-composition method has been used to investigate the influence of magnesium ion a t constant supersaturation upon the crystallization of DCPD and OCP under conditions in which these were the only precipitating phases.
Experimental Section Reagent grade chemicals and triply distilled carbon dioxide free water were used, and crystallization experiments were made in a nitrogen atmosphere maintained by bubbling with the presaturated gas. Phosphate standards were prepared from potassium dihydrogen phosphate (J.T. Baker Co., Ultrex), which was dried a t 105 OC. Calcium nitrate solutions were prepared from the tetrahydrate (Alfa Products, Ultrapure) and magnesium nitrate solutions from the hexahydrate (Baker Analyzed Reagent). Calcium and phosphate analyses were made by using a n atomic absorption spectrometer (Perkin-Elmer Model 503) for calcium and magnesium (*0.3%) and a spectrophotometer (Cary Model 210) for phosphate as the vanadomolybdate (*0.1%).19~pSpecific surface areas (SSA) were measured by nitrogen adsorption with 30/70 nitrogen/helium mixture (Quantasorb 11,Quantachrome, Greenvale, NY). DCPD seed material was prepared as previously describedz1and OCP crystals (molar calcium/phosphate ratio 1.336 i 0.01) were obtained by a constant-composition crystallization method. After aging, both materials were filtered and dried to constant weight in a desiccator over silica gel and used in subsequent experiments as seed material. The phases were characterized by X-ray diffraction (Phillips XRG 3,000 X-ray difractometer Cu K a radiation) and infrared spectroscopy (Perkin-Elmer Model 457). In a typical constant solution composition experiment, stable (8) Hayek, E.; Newesely, H. Monatsh. Chem. 1958, 89, 88. (9) Rowles, S. L. Bull. SOC. Chim. Fr. 1958, 1797. (10) Eanes,E.D.; Gillesen, I. H.; Posner, A. S. Nature (London)1965, 208, 365. (11) Eanes, E. D.; Meyer, J. L. Calcif. Tissue Res. 1977, 23, 259. (12) Davies, C. W.; Jones, A. L. Faraday SOC.Discuss. 1949, 5, 103. (13) Nancollas, G . H.; Purdie, N. Quart. Reu., Chem. SOC. 1964,18,1. (14) Nancollas, G. H.; Mohan, M. S. Arch. Oral B i d . 1970, 15, 731. (15) Tomson, M. B.; Nancollas, G. H. Science (Washington, D.C.) 1978,200, 1059. (16) Koutsoukos, P. G.; Amjad, Z.; Tomson, M. B.; Nancollas, G. H. J . Am. Chem. SOC.1980, 102, 1553. (17) Hohl, H.; Koutsoukos, P. G.; Nancollas, G. H. J. Cryst. Growth 1983, 57, 325. (18) Heughebaert, J. C.; Nancollas, G. H. J . Phys. Chem. 1984, 88, 2478. (19) Gee, A.; Deitz, V. R. Anal. Chem. 1953, 25, 1320. (20) Tomson, M. B.; Barone, J. P.; Nancollas, G. H. A t . Absorpt. Newsl. 1977, 16, 1179. (21) Marshall, R. W.; Nancollas, G . H. J. Phys. Chem. 1969, 73, 3838.
Salimi, Heughebaert, a n d Nancollas
Table I. Thermodynamic Solubility Products phase K., (mol L-lP ref DCPD 1.87 x 10-7" 25 OCP 5.0 x 10-500 26 TCP 2.83 x 10-300 27 HAP 2.35 x 10-590 28 MgHP04.3H20 1.5 X 29 Mg&'04)~-22HzO 8.0 x 1044b 29 Mg&'04)2*8HzO 6.3 X 29 ~~
"At 37 "C. b A t 25 "C. supersaturated solutions were prepared in a water-jacketed Pyrex cell by mixing solutions containing calcium, magnesium, and phosphate ions together with sufficient potassium nitrate to achieve the desired ionic strength, the p H was adjusted with dilute potassium hydroxide (Dilut-it, J.T. Baker Co.). Following verification of solution stability, a known amount of seed crystals was added to induce precipitation. The onset of growth resulted in a lowering of p H of the solution, which was immediately restored to its preset value by the simultaneous addition of two titrant solutions containing (1)calcium magnesium potassium nitrate and (2) potassium phosphate potassium hydroxide, from a pH-stat16 (Combititrator Model 3D Metrohm Brinkmann Instruments, Westbury, NY). The titrant concentrations were calculated to give a molar calcium/phosphate ratio of 1.00 and 1.33 for DCPD and OCP growth, respectively, after correction for dilution. During the experiments, the potassium nitrate present maintained the ionic strength constant to within *0.2%. Since the results of preliminary experiments, in which the concentration of magnesium ions was followed by atomic absorption, showed that they were not incorporated into the growing crystalline phases, the tritrant solutions contained sufficient magnesium nitrate solution to compensate only for the dilution effect. The pH was measured with a glass silver/silver chloride pair which was standardized before and after each experiment with National Bureau of Standard buffer solutions at p H 4.03 and 6.84.22 Titrant addition was continuously monitored, and in addition, aliquots of slurry were periodically withdrawn from the crystallization cell and filtered ( 0 . 2 2 - ~ m filters, Millipore, Bedford, MA) and the solution phase was analyzed for calcium and phosphate. The solid phases were dried to constant weight in a desiccator containing silica gel and were examined by X-ray diffraction, infrared spectroscopy, and scanning electron microscopy (IS1Model 11).
+
+
+
Results and Discussion Concentrations of ionic species were calculated from previously reported dissociation constants for phosphoric acid and equilibrium constants for the formation of ion pairs involving calcium, magnesium, and phosphate ions;lE for the magnesium species, the association constants K(MgHP04)and K(MgOH+) were 741 L mol-' 23 and 380 L respectively. An iterative computational procedure for the ionic strength, involving mass balance and electroneutrality expressions, was used. In the experiments containing magnesium ions, it was important to maintain the ionic strength constant by adjusting the concentration of added inert electrolyte (potassium nitrate). The driving force for crystallization was expressed as a free energy of transfer, AG, from supersaturated to saturated solution: AG = -RT In ( T / K ~ ~ ) ~ / ~ (22) Bates, R. G. J. Res. Nat. Bur. Stand., Sect. A 1962, 66, 179. (23) Tabor, H.; Hastings, A. B. J. B i d . Chem. 1943, 148, 627. (24) Stock, D. I.; Davies, C. W. Trans. Faraday SOC. 1948, 44, 856. (25) Marshall, R. W. Ph. D. Thesis, SUNYAB, Buffalo, 1970. (26) Shyu, L. J.; Perez, L.; Zawacki, S. J.; Heughebaert, J. C.; Nancollas, G . H. J. Dental Res. 1983, 62, 398. (27) Gregory, T. M.; Moreno, E. C.; Patel, J. M.; Brown, W. E. J . Res. Nut. Bur. Stand., Sect A 1974, 78, 667. (28) McDowell, H.; Gregory, T. M.; Brown, W. E. J. Res. Not. Bur. Stand., Sect A 1977, 81, 273. (29) Gurney; Smith, J. P. Trans. Faraday SOC. 1963,59, 1585.
Langmuir, Vol. 1, No. 1, 1985 121
Crystal Growth of Calcium Phosphates
Table 11. Growth of DCPD Seed Crystals in the Presence of Magnesium Ions at 37 "Ca initial conditions,b mol L-' expt 51 46 54 187 52 188 55 189 190
1 0 3 ~ ~ ~ 1 0 3 ~ ~ ~ 103[KN03]
0.0 0.160 0.320 2.10 3.00 4.00 5.00 7.00 10.0
9.35 9.35 9.36 9.39 9.40 9.42 9.43 9.47 9.51
AG, kJ mol-l MgHP04.3H20 Mg3(P04)2.8H20
113.6 113.1 112.7 107.3 104.6 101.6 98.6 92.6 83.6
6.9 6.0 3.6 3.2 2.8 2.5 2.1 1.6
1o4(rate),mol DCP min-' m-2 1.68 1.44 1.54 1.73 1.76 1.73 1.64 1.70 1.59
10.3
9.2 5.3 5.8 5.3 5.0 4.5 3.9
a Tc. and Tm are total molar concentrations of calcium and magnesium ions, respectively. Tc, = Tphaphats = 1.0. bpH 5.55, ionic strength 0.150 mol L-', reaction volume 250 mL, -35 mg of DCPD seed (SSA = 1.1m2 g-I). AG(DCPD) = -0.92 kJ mol-', AG(0CP) = -1.51 kJ mol-', AG(TCP) = -1.05 kJ mol-', AG(HAP) = -4.14 kJ mol-'.
expt
Table 111. Growth of OCP Crystals in the Presence of Magnesium Ions at 37 "C initial conditions," mol L-' AG, kJ mol-' 1O6(rate),mol MgHP04.3H20 Mg3(P04)2.8H20 OCP min-' m-2 1 0 3 ~ ~ ~ 1 0 3 ~ ~ 103[KN03]
183b 180b 17gb 178b 176* 182b 174* 181b 191c 191c 192c 193c 194c
4.31 4.32 4.33 4.34 4.35 4.36 4.40 4.44 3.704 3.732 3.767 3.799 3.845
83.6 82.3 81.0 79.7 78.4 75.8 70.4 63.6 85.9 80.9 74.6 68.8 60.3
0
0.430 0.870 1.30 1.74 2.62 4.40 6.66 0 1.76 3.77 5.70 8.50
5.6 4.7 4.1 3.8 3.3 2.6 2.1
7.6 6.5 5.9 5.4 4.8 4.1 3.4
4.00 2.97 2.50 1.96
5.62 4.38 3.81 2.92
1.66 1.50 1.08 1.04 0.99 0.76 0.68 0.23 0.234 0.152 0.093 0.075 0.064
'pH 6.00, ionic strength 0.100 mol L-l, reaction volume 250 mL, -35 mg of OCP seed (SSA = 8.5 m2 g-'), Tc,:Tphosphate = 1.33. bAG(DCPD) = -0.06 kJ mol-', AG(0CP) = -1.34 kJ mol-', AG(TCP) = -1.12 kJ mol-l, AG(HAP) = -4.41 kJ mol-'. CAG(DCPD)= -0.32 kJ mol-', AG(0CP) = -1.02 kJ mol-', AG(TCP) = -0.75 kJ mol-', AG(HAP) = -4.09 kJ mol-'.
/
92,
/
/
I432
0
v . 20
40
60
80
100
Time f min
120
140
I60
I
i~me(m(n1
Figure 1. Constanbcompitiongrowth of DCPD. Recorder trace of titrant volume against time. Experiment 55,32.1 mg of DCPD seed crystals. ( 0 )Growth curve corrected for change in surface area; G%, growth extent with respect to original seed crystals. where ?r is the activity product, K, is the thermodynamic solubility product for the calcium phosphate phase precipitated, and Y is the number of ions in the formula unit. The solubility data used in these calculations are summarized in Table I. The experimental conditions for the DCPD and OCP reactions are summarized in Tables I1 and I11 , respectively. A typical plot of titrant volume needed to maintain the calcium and phosphate activities constant during DCPD seeded crystallization in experiment 55 is given in Figure 1. Changes in surface area were calculated from these data, assuming the crystallites to be perfect spheres,
Figure 2. Constant-composition growth of OCP. Plots of titrant volume against time, experiment 183 ( O ) , experiment 174 ( 0 ) . Curves corrected for surface area change; experiment 183 ( o ) , experiment 174 (W).
and the ordinate values, normalized to initial seed surface area, are also shown as corrected curves in Figure 1. The linearity of these growth curves is striking and is observed for extents of growth of more than 500% with respect to the original seed crystals. The slopes of the lines are used to calculate the growth rates expressed as mol of DCPD/m2 of surface in Table 11. Typical growth plots for OCP are presented in Figure 2, and the corrected linear plots were again obtained by assuming a simple mas^)^/^ relationship to express the total seed surface area. Typical analytical data during the reactions are summarized in Table IV, and it can be seen that the calcium, phosphate, and magnesium ion concentrations remained
122 Langmuir, Vol. I, No. I , 1985
Salimi, Heughebaert, and Nancollas R R,- R
Table IV. Crystal Growth in the Presence of Magnesium Ions 1 0 ~ 7 7 ~10377phasphate, ~~ 10~77,; 9i expt mol L-* mol L-' mol L' pH seed growth" 88 7.52 7.52 10.00 5.60 DCPD 0 88 7.56 7.57 10.01 5.60 DCPD 250 126 4.50 3.37 0.495 6.00 OCP 0 4.50 3.42 0.496 6.00 OCP 126 56
"Extent of reaction expressed as percentage of growth relative to the amount of initial seed.
,[Ma]-', M-' 0
500
1000
1500
Figure 4. Influence of magnesium ions on the rate of OCP crystallization. Rate data plotted according to a Langmiur-type equation. The bars are calculated on the basis of a 5 % uncertainty in the measured rates of crystallization.
U
2
4
6
8
10 [Mgl.mM
Figure 3. Normalized rates of growth of DCPD, OCP, and HAP in the presence of magnesium ions. R and Ro are the rates of growth (mol m i d m-2) in the presence and absence of magnesium, respectively. AG(0CP) = -1.34 kJ mol-' (0) and -1.02 kJ mol-' ( 0 ) . HAP data from ref 31.
constant even during the long-term experiments (2 and 5 h for experiments 88 and 126, respectively). Moreover, even in the presence of relatively high concentrations of magnesium ion, no evidence for magnesium ion incorporation into the developing solid phases could be observed. It can be seen in Figure 3 that the presence of magnesium ions has no detectable effect on the rate of DCPD crystallization, and the same result was true a t lower supersaturation, AG(DCPD) = -0.70 kJ mol-'. In contrast, magnesium ion markedly retards the rate of seeded OCP crystallization, to almost the same extent a t both supersaturations studied, AG(0CP) = -1.02 and -1.34 kJ mol-' (Table 111). In these experiments, it should be noted that although the solutions were supersaturated with respect to TCP and HAP, the formation of TCP in the presence of magnesium ion could not be detected, despite the suggestion that magnesium ions may stabilize TCP-like structure^.^^ However, these previous studies involved spontaneous precipitation experiments in which magnesium ions could be incorporated a t the nucleation stage. The retardation experiments using OCP seed crystals were all made a t constant ionic strength and constant supersaturation with respect to OCP. Therefore the decrease in the rate of precipitation of OCP in the presence of magnesium ions could be attributed to their adsorption at active growth sites on the OCP crystal surfaces. As(30)LeGeros, R. Z.; LeGeros, J. P. In "Phosphate Minerals"; Nriagu, J. O.,Moore, P. B., Eds.; Springer-Verlag: Berlin, 1984; p 351.
suming a steady-state adsorption/desorption and the absence of interaction between growth sites, a Langmuir-type adsorption isotherm can be tested by plotting Ro/(Ro- R) against [Mgl-' where R and Ro are the rates of OCP crystallization in the presence and absence of magnesium ions, respectively. The linearity of the resulting plot in Figure 4 suggests that magnesium ions retard the rate of precipitation by adsorption at active growth sites on the crystal surfaces. It is interesting to compare the influence of magnesium ion on the rates of crystallization of DCPD, OCP, and HAP, normalized with respect to the initial rates of growth, R/Ro. These data are plotted in Figure 3 as a function of magnesium ion concentration. It can be seen that magnesium ions have a marked inhibiting effect on HAP growth, a lesser effect on OCP, and almost no influence on the growth of DCPD. The data suggest that in biological systems, where the concentration of magnesium ion may be appreciable, their specific effect in inhibiting HAP and the marked reduction in the rate of OCP crystallization may stabilize the more acidic hydrated calcium phosphates such as DCPD during the precipitation reactions. The numerous observations that brushite is a viable phase during biological mineralization may well be traced to the presence of appreciable concentrations of magnesium ions in these systems. The Langmuir-type adsorption equation was also used to interpret the inhibition of HAP mineralization in the presence of magnesium ions.31 On the basis of this model, the effective adsorption affinities calculated from the slopes of the lines such as that shown in Figure 4 are 5.5 X lo2 and 3.7 X lo5 L mol-' for OCP and HAP, respectively. Acknowledgment. This work was supported, in part, by grants from the Stauffer Chemical Co. and the National Institutes of Dental Research of the National Institutes of Health (Grant DE03223). Registry No. DCPD, 7789-77-7; OCP, 14096-86-7; Mg, 7439-95-4. (31) Amjad, Z.; Koutsoukos, P. G.; Nancollas, G. H. J. ColZoid InterScr., in press.
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