The Growth of Calcium Phosphates - American Chemical Society

32 The Growth of Calcium Phosphates 1

2

1,3

1

S. J. Zawacki, P. B. Koutsoukos, M. H. Salimi , and G. H. Nancollas State University of New York at Buffalo, Department of Chemistry, Buffalo, NY 14214 University of Patras, Department of Chemistry, Physical Chemistry Laboratory, Patras, Greece

1

Downloaded by CORNELL UNIV on June 3, 2017 | http://pubs.acs.org Publication Date: November 13, 1987 | doi: 10.1021/bk-1987-0323.ch032

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The interactions of ions with growing calcium phosphate interfaces were investigated. The effect of various background electrolytes such as NaCl, KCl, and KNO on the growth rate of hydroxyapatite HAP was measured, using a constant solution composition method. Parallel electrophoretic mobility measurements have also been made. The growth rate of HAP is markedly inhibited in the presence of magnesium ions which also induce a reversal of surface charge. The rate of octacalcium phosphate (OCP) crystallization is reduced in the presence of magnesium ion, although to a lesser extent than HAP. In contrast, dicalcium phosphate dihydrate (DCPD) crystallization is uneffected by magnesium. Strontium ions reduce the growth of both HAP and DCPD, and, unlike magnesium, are incorporated into these grown phases. 3

The growth of calcium phosphate salts is of importance both in the environment and in biological mineralization. In recent years there has been a resurgence of interest in these minerals because of t h e i r involvement in areas such as the removal of phosphate from waste water, the fate of elements such as aluminum, i r o n , and other heavy metals in the formation of lake and ocean sediments, and in industry where t h e production of scale on metal surfaces is a continuing problem. The increase in phosphate concentrations in lakes and rivers near heavily populated areas is another reason why the elucidation of the mechanism of p r e c i p i t a t i o n and the nature of t h e phases which form are problems of considerable importance (1,2). In the environment, the adsorption of metal ions on the surface of calcium phosphate salts may serve t o immobilize them in natural waters. In this process, changes in morphology and stoichiometry of the calcium phosphate crystals may prevent them from forming hard scale deposits when these waters are used industrially in applications such as cooling towers. Since both calcium and phosphate concentrations may be r e l a t i v e l y high, perhaps through the use of lime additions f o r the removal of phosphates from sewage, calcium phosphate precipitation may be o f p a r t i c u l a r importance. Higher phosphate levels are also being encountered in cooling waters due to increased water r e - u s e , the use of lower quality sewage p l a n t e f f l u e n t and c o r r o s i o n i n h i b i t o r s which a r e degraded t o o r t h o p h o s p h a t e . J

Current address: Department of Environmental Science and Engineering, Rice University, P O Box 1892, Houston, T X 77251. 0097-6156/ 86/ 0323-0650S06.00/ 0 © 1986 American Chemical Society

Davis and Hayes; Geochemical Processes at Mineral Surfaces ACS Symposium Series; American Chemical Society: Washington, DC, 1987.

32.

ZAWACKI E T AL.

Growth

of Calcium

Phosphates

651

Calcium phosphate precipitation may also be involved in the f i x a t i o n of phosphate f e r t i l i z e r in soils. Studies of the uptake of phosphate on calcium carbonate surfaces at low phosphate concentrations t y p i c a l of those in soils, r e v e a l t h a t the threshold concentration for the p r e c i p i t a t i o n of the calcium phosphate phases from solution is considerably increased in the pH range 8.5 9.0 (3). It was concluded t h a t the presence of carbonate ion from the c a l c i t e inhibits the nucleation of calcium phosphate phases under these conditions. A recent study of the seeded c r y s t a l growth of c a l c i t e from metastable supersaturated solutions of calcium carbonate, has shown t h a t the presence of orthophosphate ion at a concentration as low as 1 0 " mol L " and a pH of 8.5 has a remarkable inhibiting influence on the r a t e of c r y s t a l l i z a t i o n (4). A seeded growth study of the influence of carbonate on hydroxyapatite c r y s t a l l i z a t i o n has also shown an appreciable inhibiting influence of carbonate ion.(5). Despite the importance of the precipitation of calcium phosphates, there is s t i l l considerable uncertainty as to the nature of the phases formed in the early stages of the precipitation reactions under differing conditions of supersaturation, pH, and temperature. Although thermodynamic considerations y i e l d the driving f o r c e for the p r e c i p i t a t i o n , the course of the r e a c t i o n is frequently mediated by kinetic f a c t o r s . Whether dicalcium phosphate dihydrate (CaHPO^HoO, DCPD), octacalcium phosphate (Ca H(P0 i)3, · 2°> ) > hydroxyapatite ( C a ( P 0 ) ( 0 H ) , HAP), amorphous calcium phosphate (ACP), or a d e f e c t apatite form from aqueous solution depends both upon the driving f o r c e f o r the p r e c i p i t a t i o n and upon the i n i t i a t i n g surface phase. Thermodynamically, the r e l a t i v e supersaturation, o*> is given by


6

4

Z

2

5 H

0 C P

5

σ=

4

( Ι Ρ ^ - Κ ^ ) / ^

1

1

3

7

*

where ν represents the number of ions in a formula unit of a calcium phosphate. The ionic a c t i v i t y product, IP, of the solution with respect to the three most dominant calcium phosphate phases are: DCPD : IP = ( C a ) ( H P 0 | ) , ν = 2 2+

OCP : IP = ( C a

2 +

) (HP0^)(P0 4) , ν = 7 4

3

2

HAP : IP = ( C a ) ( P 0 4 ) ( 0 H ) , ν = 9 2 +

5

3

3

K , the thermodynamic solubility values at 37°C are, 1.87 χ 1 0 " ( m o l L " ) 7

s o

1

2

for

DCPD, 5.0 χ 1 0 " ( m o l L " ) f o r OCP, and 2.35 x 1 0 ' (mol L " ) for apatite. The driving f o r c e for c r y s t a l l i z a t i o n is expressed as a free energy of t r a n s f e r , AG, of an average ion of the calcium phosphate from supersaturated to a saturated solution: 5 0

1

7

AG

5 9

1

9

= -RT ΙπίΙΡ/Κ^) ^ 1

The experimental conditions, f r e e energies, and the observed rates of growth of the different calcium phosphates at pH 6.0 and at 37° C are summarized in Table 1. Although thermodynamically, HAP may be the preferred phase, k i n e t i c a l l y i t has a slow growth rate even though i t has the highest thermodynamic driving f o r c e and i t is quite sensitive to the presence of other ions in the supernatant solution. Moreover, more acidic precursor phases may persist for long periods, especially as surface components, without conversion to the thermodynamically most stable phase.



5.20

4.60

4.30

4.50

4.10

3.60

5.05

3.12

1.40

DCPD

DCPD

OCP

OCP

OCP

HAP

HAP

HAP

1

0.84

1.87

3.03

2.70

3.08

3.38

4.30

4.60

5.20

Solution Concentration /10-3 mol L" Calcium Phosphate

DCPD

Phase

1.00

3.00

5.00

0.45

0.61

0.74

0.17

0.24

0.39

phase

5.93

1.93

-0.42

0.38

0.62

-0.16

-0.46

-0.56

-0.85

1

AG(DCPD) /KJ mol"

1

+1.46

-0.54

-1.71

-0.96

-1.23

-1.43

-1.59

-1.73

-1;98

1

AG(OCP) , /KJ mol"

-1.78

-3.58

-4.62

-4,02

-4.30

-4.50

-4.62

-4.77

-5.03

1

AG(HAP) τ /KJ mol"

0.100 mol L' KNO3 background e l e c t r o l y t e

Experimental rates of growth, pH 6.00, 37°C,

Table 1


2.43 χ 10"

5.67 χ 10"

2.68 χ 10*

0.23 χ 10"

0.63 χ 10"

6

1.94 χ TO"

0.61 χ 10~

0.93 χ 10"

4

9

26

26

26 8

6

6

6

6

6

6

Ref

7

6

4

4

?

6

3.32 χ 10"

Growth Rate /mol min"*' m

32.

ZAWACKI E T A L .

Growth

of Calcium

Modifications of surface layers due to l a t t i c e substitution or adsorption of other ions present in solution may change the course of the reactions taking place at the solid/liquid interface even though the uptake may be undetectable by normal solution analytical techniques. Thus i t has been shown by electrophoretic mobility measurements, (6,7) t h a t suspension of synthetic HAP in a solution saturated with respect to caTcite displaces the i s o e l e c t r i c point almost 3 pH units to the value (pH = 10) found for c a l c i t e c r y s t a l l i t e s . In p r a c t i c e , therefore, the presence of " i n e r t " ions may markedly influence the behavior of precipitated minerals with respect to t h e i r rates of c r y s t a l l i z a t i o n , adsorption of foreign ions, and electrokinetic properties. In the environment, the presence of other alkaline earth cations such as magnesium and strontium may also markedly influence the course of the calcium phosphate precipitation. In natural j ^ a t e r svstems, magnesium ion concentrations may be as high as 5 χ 1 0 " mol L " (1), while in biological c a l c i f i c a t i o n , magnesium concentrations ranging from 0.5% in outer tooth enamel layers to 2% in the innermost dentine are l i k e l y to have important consequences on the rate of remineralization of carious enamel (8). It has been suggested t h a t magnesium ions k i n e t i c a l l y hinder the nucTeation and subsequent growth of HAP by competing for l a t t i c e sites with the chemically similar but larger calcium ions (9). It was shown t h a t in the presence of magnesium ions, magnesium-containing t r i c a l c i u m phosphate was formed (10,11). In contrast to the influence of magnesium ion which is v i r t u a l l y exc1ïï3ed from the growing calcium phosphates, strontium is readily incorporated into HAP l a t t i c e s because of the s i m i l a r i t y of i t s ionic radius with that of calcium (12,13). Thus a complete series of solid solutions can be prepared with l a t t i c e parameters linearly dependent upon the extent of strontium incorporation into the apatite l a t t i c e (14,15). The interest in the incorporation of strontium into calcium phosphate stems from the concern about the S r content of bones and t e e t h . Moreover, i t has been suggested t h a t the lack of strontium in the diet causes a high incidence of dental caries and poor growth conditions (16). In the present work, a constant composition method has been used to investigate the growth of HAP from solutions of low supersaturation and in the presence of different background electrolytes. The influence of magnesium and strontium ions both on the rate of c r y s t a l l i z a t i o n and upon the e l e c t r o k i n e t i c properties of the c r y s t a l l i t e surfaces has also been investigated. 2


653

Phosphates

r

9 O

Experimental Experiments were made in a nitrogen atmosphere using reagent grade chemicals and t r i p l y distilled carbon dioxide-free water. Standard solutions of phosphate were prepared from potassium dihydrogen phosphate (J.T. Baker Co., Ultrex grade), after drying at 105 C. Ultrapure calcium nitrate tetrahydrate (ATfa Products) and reagent grade magnesium and strontium nitrates (J.T. Baker Co.) were used to prepare standard solutions. Alkaline earths were determined by atomic absorption spectrophotometry (Perkin Elmer Model 503); phosphate was determined spectrophoto metrically as the phosophovanodomolybdate complex as described previously (17). A l l standard solutions were prepared using the same electrolyte composition as the growth media. Specific surface area (SSA) was measured by BET nitrogen adsorption (30/70 nitrogen/helium mixture, Quantasorb II, Quantachrome, Greenvale, N.Y.). HAP seed crystals were prepared by the method of Nancollas and Mohan (18) using calcium nitrate and potassium dihydrogen phosphate with potassium hydroxide for the control of pH. Crystal composition was verified by infrared spectroscopy (Perkin Elmer grating infrared spectrometer e



654

G E O C H E M I C A L P R O C E S S E S AT M I N E R A L S U R F A C E S

Model 467), x - r a y powder diffracton (Philips XRG-3000, x - r a y diffractometer, CuKoc radiation Ni f i l t e r ) , and scanning electron microscopy, SEM (ISI scanning electron microscope, Model II). Infrared spectra and powder diffraction data were in agreement with the published values for apatite (19,20). Chemical aaalysis of the solid gave a molar r a t i o of Ca/P = 1.64 +0.01; the SSA was 21.5 m g" · Crystallization experiments were made at 37°C in a water thermostatted double-walled vessel in metastable supersaturated solutions using the constant composition method. Presaturated nitrogen gas, at 37°C was bubbled through the solutions which were stirred with Teflon coated magnetic stirring bars. The s t a b i l i t y of supersaturated solutions prepared by mixing calcium n i t r a t e , potassium phosphate, and potassium hydroxide solutions was verified by the constancy of pH for a period of at least 4h. Following the introduction of seed crystals, c r y s t a l l i z a t i o n started immediately, and two t i t r a n t solutions consisting of (i) calcium and potassium n i t r a t e , and (ii) potassium phosphate and potassium hydroxide were added automatically from mechanically coupled burets (Metrohm Herisau, Model 3D Combititrator) in order to maintain the pH constant. The pH was measured by means of a glass electrode together with a silver/silver chloride reference electrode separated from the c e l l solution by means of an intermediate salt-bridge consisting of 0.1 mol L " * potassium nitrate. C r y s t a l l i z a t i o n experiments in the presence of magnesium and strontium ions required a third t i t r a n t containing these ions in order to compensate f o r dilution effects during the c r y s t a l l i z a t i o n reactions. Titrant addition was continuously monitored and aliquots were periodically withdrawn from the precipitation c e l l , f i l t e r e d (0.2 jim f i l t e r s , Millipore, Bedford, Ma.), and the solution phase was analyzed for divalent metal and phosphate ions, to verify constancy of composition. Typical constant composition r a t e curves for the dominant calcium phosphate phases at pH = 6.0 are shown in Figure 1. The adsorption of magnesium by the HAP substrates was investigated by equilibrating samples of the solid with solutions of calcium phosphate in 0.1 mol L " potassium nitrate calculated to be saturated with respect to HAP and with pH adjusted to 8.0 +0.1. The polycarbonate equilibration vials were gently rotated end over end for 24h at 37*C. Following this period, solid and liquid phases were separated by c e n t r i f ugation and the supernatant analyzed for magnesium. The electrophoretic mobilities of the HAP particles in the presence and absence of adsorbed magnesium were measured using a Rank Microelectrophoresis Mark II instrument with a four electrode c y l i n d r i c a l c e l l . The HAP solutions containing p r e - c a l c u l a t e d saturation concentrations of calcium and phosphate were equilibrated with the solid phases overnight and the pH was adjusted to the required values by the addition of potassium hydroxide or n i t r i c acid. Mobility measurements were made on at least f o r t y particles. 1

Results and Discussions In order to prepare calcium phosphate solutions of known supersaturation with respect to each of the phases, i t was necessary to calculate the a c t i v i t i e s of the free ionic species by successive approximation for the ionic strength, from phosphate protonation, calcium phosphate, magnesium phosphate, and strontium phosphate i o n - p a i r equilibrium constants together with mass balance and e l e c t r o n e u t r a l i t y expressions as described previously (8,21). In the experiments containing magnesium or strontium ions, i t was important to maintain the ionic strength constant by adjusting the concentration of added background electrolyte (potassium n i t r a t e ) .


32.

ZAWACKI ET A L .

Growth

of Calcium

655

Phosphates

In a conventional study of the seeded growth of HAP crystals, in which the calcium and phosphate concentrations were allowed to decrease during the reactions, i t was shown t h a t the rate of c r y s t a l l i z a t i o n varied depending upon the nature of t h e background e l e c t r o l y t e (22). Towards t h e end of t h e growth r e a c t i o n (1,000 min) the extent of c r y s t a l growth increased in the order K C l < CsCl < NH C1 < L i C l < NaCl at pH = 7.4. Moreover, i t was shown t h a t the precipitated solids at these extended times of reaction contained from 5.5 to 6.0 mol % of sodium i n the presence of sodium chloride but almost no potassium i n t h e presence of potassium chloride. These results were i n general agreement with those of Newman and co-workers (23), who found t h a t sodium could replace calcium in the calcium phosphate solid whereas potassium was reversibly adsorbed on t h e surface. A disadvantage o f the conventional precipitation method in which the supersaturation was allowed t o decrease during t h e reactions, was t h a t different calcium phosphate phases could form and subsequently dissolve during the course of the reactions. In the present work, the constant composition method was used t o investigate t h e influence o f sodium chloride, potassium chloride, and potassium n i t r a t e , as background e l e c t r o l y t e upon t h e rate of c r y s t a l l i z a t i o n of HAP i n solutions supersaturated only with respect t o this phase. These experiments were made in solutions containing t o t a l concentrations of c a l c i u m , T ç , and phosphate, T , of 1.0 χ 1 0 " ^ and 0.6 χ 1 0 ~


4

a

3

p

mol L ~ * , respectively, pH = 7.0, 3 7 T , with ionic strength made up t o 0.10 mol L~* using background e l e c t r o l y t e and 50.5 mg o f inoculating apatite seed (21.5 m g" ). The rates of c r y s t a l l i z a t i o n (+5%) i n NaCl, K C l and KNO3, background electrolytes were 2.34 χ 1 0 " , 2.09 χ 1 0 " and 1.92 χ 1 0 " mol apatite m~ min , respectively. To interpret the order of the crystallization rate NaCl > K C l > KNOq, electrophoretic mobility measurements were made on HAP particles suspended i n solutions of compositions similar to those of the c r y s t a l l i z a t i o n experiments. In addition, potentiometric t i t r a t i o n s were made in the presence o f these electrolytes (24) in order to determine the point of zero charge, pzc. Electrophoretic mobility results as a function of pH, shown i n figure 2, r e v e a l t h a t i n the case of potassium nitrate no specific i n t e r a c t i o n of t h e e l e c t r o l y t e with HAP takes place since the i s o e l e c t r i c point (iep), pH = 6.7 in figure 2 coincides with t h a t corresponding to the pzc. In potassium nitrate solutions, the surface charge of HAP as determined by microelectrophoresis did not show any change f o r extents of growth up t o 60% of new material deposited on the inoculating seed. In contrast, in the presence of potassium chloride, the iep (pH = 6.4 in figure 2) is markedly different from the pzc value (at pH = 8.5) indicating a stronger interaction of C T compared t o NO3 i o n . In t h e case of sodium chloride (figure 2) the iep is shifted to pH = 5.20 reflecting the combined e f f e c t o f both sodium and chloride ions. Under the conditions of t h e c r y s t a l l i z a t i o n experiments, i t can be seen from figure 2 t h a t at pH = 7, the HAP surface carries the most negative charge in the presence of sodium chloride. The observed decreasing negative charge, NaCl > K C l > KNO3 follows the same trend as the corresponding rates of c r y s t a l l i z a t i o n . 2

1

8

8

8

2

;

An advantage of the constant composition technique is t h a t r e l a t i v e l y large extents o f growth and enhanced c r y s t a l l i n i t y can be achieved at low supersaturations. Improved crystallinity of the particles during c r y s t a l l i z a t i o n is r e f l e c t e d in lower specific surface areas of the solid phases; x - r a y powder diffractograms of the solid phases removed from the c r y s t a l l i z a t i o n c e l l also show increases in sharpness. Experiments i n which crystal growth was allowed to proceed until f i v e or six times the amount of


GEOCHEMICAL

P R O C E S S E S AT M I N E R A L S U R F A C E S


656

Figure 2. The e l e c t r o p h o r e t i c m o b i l i t y o f HAP p l o t t e d a g a i n s t pH, i n d i f f e r e n t background e l e c t r o l y t e s (0.01 mol L - l ) a t 37°C; • , NaCl ; Ο , KCI ; Ο , K N 0 . 3



32.

ZAWACKI E T A L .

Growth

of Calcium

657

Phosphates

original seed was deposited, showed striking changes in morphology in the presence of chloride ion when added as potassium chloride (25) or lithium chloride (26). In contrast to the needle-like HAP morphology grown in the absence of chloride i o n , platelike crystals were formed in its presence and measurements of the unit c e l l l a t t i c e parameters revealed a slight increase in the a axis and decrease in the c axis during c r y s t a l l i z a t i o n . Such changes may r e f l e c t the presence of chloride ion in the apatite l a t t i c e and this suggestion is supported by the electrophoretic mobility results given above. In experiments conducted in the absence of chloride ion in which the supersaturated calcium phosphate solutions and t i t r a n t solutions were prepared using calcium hydroxide and phosphoric acid, c h a r a c t e r i s t i c needle-like HAP crystallites with the required hexagonal unit c e l l l a t t i c e parameters were obtained (25). Uptake of magnesium ions by HAP surfaces at pH = 8.0 is i l l u s t r a t e d in the adsorption isotherm shown in figure 3. The adsorption markedly exceeds t h a t corresponding to a monolayer coverage (approximately 1.5/umol m " , calculated assuming t h a t the hydrated ion radius is equal to its magnesium crystal radius plus the diameter of a bound water molecule). The sharply rising adsorption at low magnesium concentrations is indicative of the high affinity between substrate and adsorbent. No plateau was observed in the isotherm and further increase of magnesium concentration in the equilibrium solution leads to magnesium uptake levels as high as 27 y mol m " (26). Studies of the influence of magnesium ion upon the c r y s t a l l i z a t i o n of calcium phosphate phases, showed l i t t l e or no evidence for incorporation of this ion into the l a t t i c e . It has been suggested t h a t the inhibition was due to adsorption of the added metal ion at the surface of the crystals (8). The results of the electrophoretic mobility measurements on HAP particles having surface concentrations of magnesium ion of 2.5 and 25 x l O " ^ μ mol m " are shown in figure 4. It can be seen t h a t adsorption of magnesium ions leads to marked changes in zeta potential. It appears t h a t magnesium uptake on HAP surfaces may not be a simple adsorption process since the i s o e l e c t r i c point is shifted in a direction opposite to t h a t expected for cation uptake. An apparent surface concentration greater than t h a t corresponding to a monolayer of hydrated magnesium ions may be attributed to p a r t i a l dehydration of the adsorbed ions. This may suggest the formation of a "surface phase" of calcium-magnesium-phosphate. The electrophoretic mobility profile as a function of pH of the HAP samples with increasing concentration of magnesium ions at the surface, approaches t h a t exhibited by Whitlockite (26). The influence of magnesium and strontium ions upon the c r y s t a l l i z a t i o n rates of calcium phosphate phases are summarized in figures 5 and 6. Previous work has shown t h a t while having no detectable e f f e c t on the growth of DCPD, magnesium ions appreciably retard the rates of OCP and HAP c r y s t a l l i z a t i o n (8). In figure 5 i t can be seen t h a t the moderate retardation of OCP c r y s t a l l i z a t i o n by magnesium ion at pH = 6.0 contrasts the much greater inhibition of HAP growth at ρ H 7.4 and ρ H 8.5 (28). Unlike magnesium, the strontium ion is readily incorporated into the growing calcium phosphate c r y s t a l l a t t i c e s {2Vj. It can be seen in figure 4 that at pH = 7.40 and ionic strength of 0.01 mol L ~ \ the inhibiting influence of magnesium ions on the c r y s t a l l i z a t i o n of HAP is considerably greater than that observed in the presence of strontium i o n . In contrast to the insensitivity of DCPD growth to the presence of magnesium i o n , small decreases in c r y s t a l l i z a t i o n rate of 6% r e l a t i v e to results in the absence of strontium accompany the incorporation of 3 mole per cent strontium into the DCPD l a t t i c e as shown in figure 6. Incorporation of 2% strontium resulted in a 40% 2

2

2


658

G E O C H E M I C A L P R O C E S S E S AT M I N E R A L S U R F A C E S


10

01 0

2

4

ι • • • I 8 10

6

C Mg ] o q

/

10"

mol

4

L

_ 1

F i g u r e _3. P l o t o f magnesium adsorbed per m o f HAP a g a i n s t the e q u i l i b r i u m magnesium c o n c e n t r a t i o n s , r e m a i n i n g i n s o l u t i o n a t pH 8 . 0 , 0.10 mol L " KN0 background e l e c t r o l y t e a t 3 7 ° C . 2

1

3

3

-2-M

·

4

1 5

1

1 6

·

P

1 7

·

1 8

·

1 9

·

h-

10

H

Figure 4. I n f l u e n c e o f magnesium a d s o r p t i o n on t h e e l e c t r o p h o r e t i c m o b i l i t y o f HAP a t 3 7 ° C , 0.01 mol L"1 KNO3 background e l e c t r o l y t e ; Ο , HAP s u r f a c e , no a d s o r p t i o n ; • , 2 . 5 /xmol m-2; , 25 jx mol m-2, Mg2+ a d s o r b e d .



32.

ZAWACKI E T A L .

Growth

0-M

of Calcium

,

1

0

,

C

M

]

1

•

4

2 /

10"

659

Phosphates

1

,

4

mol L "

,

1

6

8

μ

10

1

F i g u r e 5. I n f l u e n c e o f magnesium and s t r o n t i u m on the r a t e o f p r e c i p i t a t i o n o f a p a t i t e s a t 37°C. ftyj and R Q are the r a t e s i n the presence and absence o f metal i o n s , r e s p e c t i v e l y . • , Mg2+ on HAP, pH 8 . 5 0 , 0.01 mol L " l K C I , T = 3.0 χ 10-4 mol L - l , T = 1.8 χ 10-4 mol L - Ί ; • , Mg2+ on HAP, pH 7 . 4 0 , 0.01 mol L - l K C I , T c = 5 . 0 x 10~4 mol L - Ί , Tn = 3.0 x 10-4 mol L - l ; Ο , Sr2+ on HAP, pH 7 . 4 0 , 0.01 mol L - l KC1, T = 5.5 x 10-4 mol L - l , T = 3.3 x 10-4 l L-l; Ο , Mg2+ on OCP; pH 6.00, 0.10 mol L - l KNO3, T = 3.7 x 10-3 mol L - l , T = 2.98 x 10-3 mol L - l . C a

p

a

C a

m

p

C a

p


o


660

GEOCHEMICAL

TimQ

/

P R O C E S S E S AT M I N E R A L

SURFACES

min.

F i g u r e 6. I n f l u e n c e o f s t r o n t i u m i o n on the i n i t i a l p r e c i p i t a t i o n o f DCPD and HAP. • , DCPD, pH = 5 . 6 0 , 8 χ 1 0 " mol L " ' Ca(N03)2 and 8 χ 1 0 - 3 l Η K H P 0 , 0 . 0 7 8 mol L - l K N 0 ; • , DCPD w i t h 3% Sr i n c o r p o r a t i o n , 6.6 χ 1 0 - 3 mol H C a C l 2 + 1 . 4 χ 1 0 - 3 mol L"1 S r C l 2 and 8 x 1 0 - 3 mol L - l K H P 0 , 0 . 0 7 8 mol L - l K C l ; Ο , HAP, pH = 7 . 4 0 , 0 . 5 5 x 1 0 - 3 mol L " C a C l 2 and 0 . 3 3 mol L - l KH2PO4 0.01 mol H KCl; 0 , HAP, 2% Sr i n c o r p o r a t i o n ; φ , HAP, pH = 7 . 4 0 , 12% Sr incorporation. 3

m

o

2

4

2

3

4

1



32.

ZAWACKI E T A L .

Growth

of Calcium

Phosphates

661

retardation whereas a 12% strontium uptake reduced t h e growth r a t e by 57% (figure 5). The incorporation of strontium into the DCPD l a t t i c e was accompanied by small but significant expansions of t h e a and c l a t t i c e parameters (28) as was found f o r t h e HAP system (21). In contrast t o the growth experiments i n the presence o f magnesium i o n , the addition o f strontium markedly reduces t h e average size of the crystallites formed (21). In conclusion, i t has been found that ions frequently found i n t h e environment may have very different effects on the growth and surface properties of calcium phosphate phases. Chloride ion appears t o accelerate the reactions either due t o i t s electrostatic interactions with the apatite surface or as a r e s u l t o f substitution f o r hydroxyl ions. In contrast, sodium ions reduce the rate and may be incorporated into the growing phase (22). Strontium and magnesium ions appear t o influence t h e r a t e by very different mechanisms. The small magnesium ion has no measurable e f f e c t on t h e least thermodynamically stable and fastest growing DCPD, i t has a moderate inhibiting e f f e c t on OCP and markedly inhibits the slow growing apatite, even at micro molar concentrations. Since there is l i t t l e evidence f o r appreciable incorporation of magnesium into any of these growing c r y s t a l l i t e s , this rate reduction probably results from surface interactions. The strontium i o n , being similar in size to calcium, not only slows down the growth of DCPD and HAP, but is r e a d i l y incorporated into the growing crystals. Acknowledgments We thank the National Institute of Dental Research o f the National Institute of Health f o r a grant (Number DE03223) i n support of this work. We also acknowledge a grant by NATO (Number 614-83) to Peter G. Koutsoukos and George H. Nancollas.

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Lindsay, W.L. "Chemical Equilibria in Soils"; Wiley: New York, 1979. Brown, W.E. In "Environmental Phosphorus Handbook"; Griffith, E.J.; Beeton, A; Spencer, J.M. and Mitchel, D.T., Eds.; Wiley and Sons: New York, 1973, p.203. Boischot, P.; Coppenet, M.; Herbert, J. Ann. Agron. 1949, 19, 103. Kazmierczak, T.F. Ph.D. Thesis, SUNY, Buffalo, 1978. Koutsoukos, P.G. Ph.D. Thesis, SUNY, Buffalo, 1984. Hassan, K.A.R. MS Thesis, SUNY, Buffalo, 1984. Somasundaran, P.; Wang, Y.H.C. In "Adsorption and Surface Chemistry of Hydroxyapatite"; Misra, D.N. Ed.; Plenum Press: New York, 1980, p.129. Salimi, M.H.; Heughebaert, J.C.; Nancollas, G.H. Langmuir 1985, 1, 119. Ferguson, J.; McCarty, P.L. Environ. Sci. Technol. 1971, 5, 534. Hayek, E.; Newesely, H. Monatsch. Chem. 1958, 89, 88. Rowles, S.L. Bull. Soc. Chim. Fr. 1958, 1797. Elliot, J.C. Clin. Orthop. 1973, 93, 313. Young, R.A. Colloq. Int. C.N.R.S. No. 230 1973m 21. Colin, R.L. J. Am. Chem. Soc. 1959, 81, 5275. Hayek, E.; Petter, H. Monatsch. Chem. 1960, 91, 356. Shield, C.P.; Curzon, M.E.J.; Featherstone, J.D.B. Caries. Res. 1984, 18, 495. Tomson, M.B.; Barone, J.P.; Nancollas, G.H. At. Absorpt. Newsl. 1977, 16, 117.


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Nancollas, G.H.; Mohan, M.S. Arch. Oral Biol. 1970, 15, 731. Baddiel, K.B.; Berry, E.E. Spectrochim. Acta, 1966, 22, 1407. ASTM, "X-ray Powder Diffraction File" No. 9-77. Koutsoukos, P.G.; Nancollas, G.H. J. Phys. Chem. 1981, 85, 2403. Nancollas, G.H.; Tomazic, B. J. Phys. Chem. 1974, 78, 2218. Newman, W.F.; Toribara, T.Y.; Mulryan, B.J. Arch. Biochem. Biophys. 1962, 98, 384. Koutsoukos, P.G.; unpublished data. Koutsoukos, P.G.; Nancollas, G.H. J. Cryst. Growth 1981, 55, 369. Koutsoukos, P.G.; Zawacki, S.J.; Nancollas; G.H., in preparation. Amjad, Z.; Koutsoukos, P.G.; Nancollas, G.H. J. Coll. Int. Sci. 1984, 101, 250. Shyu, L.J. Ph.D. Thesis, SUNY, Buffalo, 1982. June 3, 1986


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