SURFACE CHEMISTRY OF BONE
Sept. 5 , 1956
t o decomposition. Heating a t temperatures below 100' and at reduced pressures proved time consuming and inadequate. Above 80' prolonged heating always resulted in the loss of some fluorine due t o hydrolysis. Successful partial dehydration was accomplished in a n evacuated desiccator over phosphorus pentoxide at 65" in approximately 50 hr. Loss of total sample weight and fluorine and phosphorus analyses indicated t h e hemihydrate formation. When held for 9 days a t these conditions, there was no additional loss in weight a n d analytical d a t a indicated pure hemihydrate. There was no particular consistency in t h e rate of dehydration a n d influencing factors were sample size, exposed surface, condition of t h e phosphorus pentoxide and degree of evacuation of the desiccator. T h e only evidence for t h e existence of a stable monohydrate was obtained by hydrating t h e hemihydrate. When placed at 20" in a constant relative humidity of 15% (approximately 3.5 mm. of water vapor) the hemihydrate gained sufficient weight t o convert theoretically t o the monohydrate. There was no additional gain in weight with time, and there was no loss in weight when placed over anhydrous calcium chloride. T h e monohydrate gained in weight t o the dihydrate value when placed in a 51% relative humidity dtmosphere a t 25". However, this second water molecule could now be removed b y drying over anhydrous calcium chloride at room temperature whereas t h e original dihydrate water could not he removed in this manner. Elevated temperatures failed t o yield t h e anhydrous salt. Decomposition with loss of fluorine resulted in every case. The theoretical composition of possible hydrates of calcium monofluorophosphate are given in Table I1 for comparison.
TABLE I1 COMPOSITION OF POSSIBLE HYDRATES Hydrate Ca, % F, % p, % CaP03F.2H20 23.0 10 9 17.8 CaP03F.H20 25.6 12.2 19.9 CaP03F.*/?HzO 27 2 12 9 21.1 CaP03F 29 0 13 8 22.5
[COSTRIBUTION F R O M
THE
DIVISIOSO F
4263 TABLE I11
SOLUBILITY OF CaPOaF.2H20IN WATER Temp.,
5 17 27
37 48 58
O C .
G. of CaPOaF/100 ml. soln.
0.486 ,476 ,417 ,390 ,438 ,486
=k 0.006
zt &
zt zt It
,006 ,006 .006 .006 .006
T h e solubility of the salt in water (Table 111) was determined b y volumetric analysis of the saturated solution and all solubilities were less than 0.5 g. of CaPOaF per 100 ml. of saturated solution. Equilibrium was reached by constant stirring of a solution containing excess solid for 24 hr. with longer periods of stirring time showing no increase in solubility. After 2 months of solid-liquid contact, the salt was found t o be negligibly soluble at room temperature in all organic solvents tested. These included: 95% ethanol, carbon tetrachloride, chloroform, 1,4-dioxane, tetrahydrofuran, ethyl acetate, furfural, pyridine, thiophene, ethylene glycol, ethylene glycol monoethyl ether and carbon disulfide. Crystallographic studies were hampered by the difficulty of obtaining large perfect crystals. However, studies revealed the dihydrate crystals had a n inclined extinction angle whose measurements varied between 33 and 37" and also a parallel extinction angle. These indicated the monoclinic crystal system for t h e dihydrate, which was also observed t o form twin crystals. T h e hemihydrate crystals appeared light brown in color and had the same general shape as the dihydrate. However, it was suspected t h a t they did not belong t o the same crystal system b u t t h a t a case of pseudomorphism had been observed. It has been noted t h a t t h e properties of monofluorophosphate compounds closely resemble those of the corresponding sulfates. This investigation shows a marked similarity between calcium monofluorophosphate and gypsum, CaSO4.2H20. NORXIAS,OKLAHOMA
DEPARTMENT O F RADIATIOSBIOLOGY, SCHOOL DENTISTRY, UNIVERSITY O F ROCHESTER]
PHARMACOLOGY,
O F ?VIEDICISE AND
The Surface Chemistry of Bone. IX. Carbonate :Phosphate Exchange1 BY W. F. KEUMAN, T. Y . TORIBAFU AND B. J. MULRYAN RECEIVEDAPRIL 25, 1956 Under carefully controlled conditions, hydroxyapatite crystals were equilibrated with bicarbonate buffers. Bicarbonat e was found t o penetrate the hydration shells of t h e crystals and, in addition, displace phosphate ions from t h e surfaces of t h e crystals. These two physicochemical processes may account for t h e large amounts of COn found in bone. T h e exchange reaction is quantitatively of t h e most physiological importance since t h e crystals of mature bone are poorly hydrated in vivo.
In 1881, Hoppe-Seyler2 attempted to describe bone salt by a formulation which included carbonate ions as part of the "molecule." Today, i t is generally recognized t h a t bone salt cannot be represented as a "molecule" but rather is best described as microcrystalline material exhibiting the The exact lattice structure of hydroxyapatite. nature of the ever-present carbonate of bone mineral is still uncertain, however, and there exists in the literature an array of suggestions and theories, none of which has been experimentally established beyond reasonable doubt.3b
None of these suggestions takes into account the recently-discovered fact t h a t crystals of hydroxyapatite, in an aqueous medium, possess hydration s h e l k 4 Therefore, a series of investigations of a model system (aqueous buffer-hydroxyapatite crystals) was conducted to clarify the problem of carbonate fixation in bone.
(1) This paper is based on work performed under contract with the United States Atomic Energy Commission a t T h e University of Rochester Atomic Energy Project, Rochester, New York. ( 2 ) F. Hoppe-Seyler, Z. Physiol. Cheni., Berlin (1881). ( 3 ) (a) S.Eisenberger, .4. Lehrman and UT. D. Turner, C h i n . Revs., 26, 257 (1940); (b) W F.Neuman and SI. W. lu'euman, ibid., 63, 1 (1933).
NAL,
Materials and Methods.--.% well-characterized preparation of crystalline hydroxyapatites-8 was used for all cquilibration studies. All other chemicals were coinmercially available, C.P. grade. T h e methods a n d apparatus used (4) W. F. Neuman, T. Y . Toribara a n d B. J. Mulryan, THISJ O U R -
76, 4239 (1953). ( 5 ) W. F. Neuman, Atomic Energy Report UR-238 (1953). (6) J . H. Weikel, J r . , W. F. S e u m a n and I. Feldman, THISJOURNAL, 76, 5202 (1954). (7) G. J . Levinskas and W. F. Neuman, J . Phrs. C h e i n . , 69, 164 (1955). (8) W . R . Stoll and W. F. Neuman, THISJ O U R N A L , 78, 1583 (1950).
for equilibration have been previously described.6 Phosphorus was analyzed b y the Fiske arid SubbaRow methodQ; calcium by either the Versene'O or flame photometric methods"; sodium by flame photometry*; and C 0 2 was analyzed gravimetrically after the gas had been released by acid and collected in ascarite tubes.
Results
portions. ii'ithin experimental error, t h e CO1-fixdtion by the solid was uninfluenced by changes in ionic strength. I n both cases, the C 0 2 in the solid phase was roughly proportional t o the concentration of bicarbonate in the solution phase. I n neither caSe was the relationship linear. These results are given in Fig. 2. T h e legend accompanying the figure presents pertinent experimental details.
Time Required for Attainment of Steady State.-Two experiments were performed t o determine the time required for apatite crystals t o equilibrate with a carbonate-containing buffer solution. I n the first experiment, 2 g. of apatite crystals was added t o 100 ml. of a n equal mixture of 1 h ' KHCOa and 1 M KCl which had been adjusted t o p H 7.5 by bubbling COSthrough the solution. -after varying time intervals of rapid stirring, the crystal suspension was centrifuged, the resultant crystalline sludge transferred t o special centrifugation cups4 and centrifuged at 10,000 X g for 2 hr. T h e hydrated crystals, freed of mechanicallyheld ~ a t e r , ~ 3were 8 then analyzed for C o r content. The averaged results of duplicate experiments are given in Fig. 1.1,
c 02 mg. /g.
IO
wet wt.
l5
t
JA -1.0 H C 0 3 - =0.5 cation = K
5t 1
IA 20
40 60 88 0 HOURS EQUILIBRATION.
-
*
15-
c 02 mg. / g. wet wt.
"1
0
u = 0.3 H C O 3 = 0.3
0.2 0.4 0.6 0.8 1.0 S o h . phase (moles/l.). Fig. 2.-A curve showing the relation between COS taken u p by the solid phase and the concentration of bicarbonate in the solution. The solid triangles represent data from experiment I employing the following experimental conditions: 1.1 = [HCOa], cation = K a + , p H 7.4, 4.5 g. apatite/ 600 ml. buffer, T = 2 5 O , 2 hr. equilibration; the open circles represent data from experiment I1 under the following conditions: p = 1.0, cation = EC-, pH 7.1, 4.5 g. apatite/ 600 buffer, T = 25", 2 hr. equilibration. I n experiment I , the extent of hl-drdtion of the crystal5 was determined by methods previously published.4 These data are given in Fig. 3 and demonstrate that, at high levels of C o r impregnation, there was a significant loss of hydration water. =\ simple calculation shows clearly t h a t only a part uf the C 0 2 found in the solid phase can be attributed to bicarbonate in the 115-dration shell'*: 0.67 ml./g. X
80
cation = N a *
,
IB
I
, . 75
2 4 6 HOURS EQUILIBRATION.
Q A
Fig. 1.--Data showing the rapidity with which COz-uptake by the solid phase attains steady-state conditions.
70
In the second experiment, conditions were altered slightly in t h a t 4.5 g. of apatite was equilibrated with 600 ml. of 0.3 31 NaHCO, a t p H 7.4. T h e hydrated crystals were isolated and analyzed a s before. These results are given in Fig. IB. U'hile no claim is made t h a t equilibrium was achieved, it is perfectly clear t h a t the crystals' uptake of CO? was very fast and remained at a constant level, within experimental error, from 5 minutes (actually it required < 2 hr. t o isolate the hydrated crystals) t o 130 hr. equilibration. Neither the nature of the cation nor the ionic strength of the buffer had any measurable effect on the time required. Accordingly, an equilibration time of 2 t o 6 hr. was adopted for all subsequent studies. Effects of Varying Bicarbonate Concentration.-Two esperiments were conducted testing the effects of varying the concentration of bicarbonate ion in the solution equilibrating with the crystals. I n the first case experiment I [HC03-1 was varied b y dissolving different quantities of NaIICOa; therefore, the ionic strength varied proportionately to the [HC0,-I. In the second case, the ionic strength vias maintained a t unity b y mixing KC1 and KHCO? in varying pro-
65
I
I
0.2
0.4 0.6 0.8 1.0 HCOB-, M. Fig. 3.--Data showing loss of hydration shell n a t e r with increasing bicarbonate concentration. Hydration water has been expressed as: wt. H20 X 100/wt. dry apatite. 3 AI HC03 = 0.67 mmole C02/g., while observed fixation was 1.05 mmole C02/g. This suggests that bicarbonate (or carbonate) ions entered iuto the crystal surface in addition t o penetrating the hydration shell. In support of this, it was found in experiment I1 t h a t COYfixation by the solid vas accompanied by a displacement of phosphate from the crystals t o the solution as seen in Fig. 4. Effect of Calcium Concentration.-It has been rcportet16 t h a t the amount of readily exchangeable phosphate ( measured with radiophosphate) associated with hydrox apatite crystals can be markedly increased b y increasing the concentration of calcium in the equilibration fluid. Accordingly, an experin~entwas performed testing whether calcium exerted n similar influence 011 c:irlmi:ite fixation. The results are given in Fig. 5 , tlic lcgcntl o f which ecint:iini calciilati
