March, 1958
377
NOTES
SURFACE CHEMISTRY OF BONE MINERAL. X. THE LACK OF INTERACTION BETWEEN SODIUM AND CARBONATE IONS'
THE
in rolutlon (mM/l.),
[K']
. , O i e " " " l l
BYW. R. STOLL AND W. F. NEUMAN From the Divieion of Pharmacolooy, Department of Radiation B~OZOQU, School o j Medicine and Dentistry, University of Rochester, Rochester, New York Received October 14, 1967
Sodium ions have been shown t o compete with calcium ions for the crystalline surface of hydroxy apatite on a mole for mole basis.2 Under physiological conditions, the magnitude of this passive sodium exchange by hydrated hydroxy apatite is adequate to account for the sodium content of bone in vivo. A similar crystal surface competition has been shown recently to exist between carbonate and phosphate ions.a In this case, although the primary ion in solution is bicarbonate, infrared studies have demonstrated that the ion in the crystal surface is the carbonate ion.4 The few studies in the literature of sodium and carbonate mobilization from bone have revealed that these two ions are closely associated, a decrease of one being associated with a decrease in the This association might be explained either on a physiological or a physicochemical basis: (a) t o deplete bone carbonate, the animal must be made acidotic and, physiologically, most acidotic states are accompanied by a sodium loss or (b) some specific interaction between sodium and carbonate ions occurs at the mineral crystal surfaces. Experiments were performed in vitro to test the physicochemical explanation. No specific physicochemical interaction was observed. Experimental Procedures Two series of equilibrating solutions were used. The first series contained varying amounts of sodium and potassium as chlorides: the second contained mixtures of sodium and potassium as bicarbonates. Each solution was 0.005 M with respect to diethylbarbituric acid which served as a buffer. The ionic strength of each soIution was approximately 0.16 M. Two grams of a sodium-free, synthetic hydroxy apatite2'8J was equilibrated with one liter of the test solut,ion. The temperature was maintained at 37' by means of a waterbath, the pH was maintained a t 7.4 & 0.1, and the time of equilibration was from 2-6 hours, a time proven adequate for the establishment of equilibrium by earlier experiments (the techniques of equilibration, of sampling the solutions, and the separation of the solid phase are described elsewhere).at* After separation from the solution, the wet solid was divided into four parts and centrifuged at a force of 8000 X gravity for 1 hour to remove the last traces of the bulk solution leaving the crystals with their hydration shells -
(1) This paper is based on work performed under contract with the United States Atomic Energy Commission a t the University of Rochester Atomic Energy Project, Roahester, New York. (2) W. R. Stoll and W. F. Neuman, J. A m . Chem. Sac., 7 8 , 1585 (1956). (3) W. F. Neuman, T. Y. Toribara, and B. J. Mulryan, ibid., 78, 4263 (1956). (4) (a) A. 9. Posner and G. Dukaerts, Ezpericntia, 10,424 (1954); (b) A. L. Underwood, T. Y. Toribara and W. F. Neuman, ibid., 77, 317 (1955). ( 5 ) W. H. Bergstrorn, J. Bid. Chem., 906, 711 (1954). (6) W.H.Bergstrom, J. Clin. Invest., 34, 997 (1955). (7) G. Nichols, Jr., and N. Nichols, Metabolism, Clin. and Exp., 6 , 438 (1956). ( 8 ) W. F. Neuman, Univ. of Rochester Atomio Energy Report. UR-238 (1953).
0
50
[No']
100
IS0
in solution (rnM./l.).
Fig. 1.-Uptake of sodium and potassium by hydrated hydroxy apatite in the presence of chlorides and bicarbonates: 0,sodium uptake from chloride solutions; 0 , sodium uptake from bicarbonate solutions; O, potassium uptake from chloride solutions; . , potassium uptake from bicarbonate solutions. intact for subsequent analysis.*sS Following centrifugation, the solid phase was dried and dissolved in a fixed quantity of dilute nitric acid. The liberated carbon dioxide was determined gravimetricalIy.3 Aliquots of the acid solutions were then taken for sodium and potassium analyses by flame photometry.2 A second aliquot was taken for chloride analyses by displacement of iodate from silver iodate by chloride. The released iodate in the supernatant was then reduced by potassium iodide and the free iodine was titrated with sodium thiosulfate, This method, an adaption of the method of Sendroylo was not affected by variations in the amount of dissolved apatite, but it was quite sensitive to variations in the concentration of acid. Samples of the equilibrating solutions were analyzed for calcium by titration with ethylenediaminetetraacetic acida and for phosphorus by the method of Fiske and SubbaRow.11 The concentration of phosphate in the hydration shell of the crystals and on the crystalline surface itself was estimated by isoto e dilution techniques using carrier-free P8'J04 as describef previousIy.l* Each set of equilibrations w&s repeated. The two sets of data were then averaged as duplicate analyses.
Results and Discussion The uptake of potassium and sodium by the hydrated hydroxy apatite is shown in Fig. 1 as a function of the sodium ion concentration in the equilibrating solution. The potassium concentration can be obtained by the relation: [K+] = 160- [Na+] in millimoles per liter. Where chloride was the anion present, the uptake of the sodium by the hydrated solid was greater than the uptake of potassium in equivalent concentrations, in excellent quantitative agreement with an earlier study.2 When bicarbonate was the anion present, the uptake of both sodium and potassium ions was increased slightly. (9) W.F. Neuman, T. Y.Toribara and B. J. Mulryan, J. A m . Cham. Soe., 7 6 , 4239 (1953).
(10) P. B. Hawk, B. L. Oser and W. H. Sumerson, "Practical Physiological Chemistry," 12th Edition, Blakiaton Co., Philadelphia, Penns., 1949,p. 885. (11) C. H.Fiske and Y. SubbaRow, J. Biol. Chem., 66, 375 (1925). (12) J. H.Weikel, W. F. Neuman and I. Feldman. J . A m . Chem. Sac., 76,5202 (1954).
NOTES
378
Vol. 62
quantity, [N5t]soln.,
(’)”a hydration shell * YK t hydration shell) .X the presence of carbonate has no specific
influence. A comparison of the effect of the anion present (chloride us. bicarbonate) on the composition of the solid phase and of the solution is given in Table I. It is true that the concentrations of sodium and of calcium ion influence the composition of the hydration shell and of the 14 However, these changes are insignificant in comparison with the changes induced by replacing chloride with bicarbonate as the anion present. In Table I, therefore, the averages of all the bicarbonate and all the chloride mixtures are given irrespective of the cation, potassium or sodium. 0
1000
2000
3000
aNa+/aCa++solution. Fig. 2.-Competition of sodium and calcium for the surface of hydroxy apatite, plotted according to mas8 law: 0,chloride solutions; 0 , bicarbonate solutions; 0 , chloride solutions from previous data.2
Previously, the conclusion was drawn2 that potassium ions do not displace calcium from the crystalline surface to any significant extent; rather, the potassium found by analysis of the apatite was assumed to be the result of a simple equilibrium between the bulk solution and the hydration shell of the solid phase. With the reasonable assumption that the content of sodium of the hydration shell is the same as the potassium content at equivalent concentrations in solutions, the difference between the total uptake of these two cations approximates the concentration of sodium actually in the crystalline surface as a result of an exchange displacement of calcium ion.2 Since sodium displaces calcium from the surface on an equal molar basis, the net sodium uptake (Nasolid - K o l i d ) should be proportional to the activity ratios of sodium to calcium in the solution.2 This function is given in Fig. 2 and, for comparison, previous data2 are also included. The same line describes the sodium-calcium relationship equally well, regardless of the anion species present in the solution. Despite the consistency of the exchange-interpretation of the earlier study12it is possible to account for the preference of the solid phase for sodium ion on the basis solely of activity coefficients. For example, the apparent activity coefficient, YK t, in the hydration shell has been calculated to be 0.55,2 suggesting the efective ionic strength of the hydration shell to be about 1.5. At this ionic strength, the activity coefficient of sodium ( Y N t~ hydration shell) can be expected to be even SlrXder than Y K + hydration shell, by about 20%.lS Thus sodium would concentrate at the crystal: solution interface and the resultant electrostatic inbalance would release other cations, principally calcium. In any case, whether sodium uptake by the solid represents true exchange or whether the ordinate in Fig. 2 represents the (13) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” Reinhold Publ. Carp., New York, N. Y., 1950,p. 562.
TABLE I EFFECT OF ANION(Cl- us. HCOI-) ON SOLID A N D SOLUTION COMPOSITION^ Compn. of solid
HzO c1
coz
Anion in s o h Chloride Ricerbonate 160 in mole/l. 152 in mole/l.
Net change
53 f 0.6 53 i 0.8 0.155 =k .005 . . . , , , , , . . . . . . . . . . . . , 0.31 f .03
Pbydration ohell
. 2 f .01 .4 f .I
Psurfaoe
.1 f
.01
.I f . 1
-0.1 f 0 . 0 2 - .3 f . 2
Compn. 02 s o h .
.24 f .006 .11 i .01 - .13f .02 .12 f .004 .41 f ,016 . 2 9 & .02 a Results are av. f av. deviation of four analyses, expressed as mmoles/g. in solid or mmoles/l. in solution. For a comparison of net changes, solid us. solution, the mmoles/g. must be multiplied by two since 2 g. of solid was equilibrated with one liter of solution. Ca
+
P
There were no detectable changes in the size of the hydration shell. Phosphate ions appeared to be displaced by carbonate ions both from the hydration shell and from the crystal surface. The average error of the total amount of phosphate displaced, 0.4 =k 0.2 mmole per gram, was too large to permit any conclusion regarding the molar ratio of the carbonate-phosphate displacement. The displaced phosphate appeared in the solution and repressed the amount of calcium dissolved from the solid. On a molar basis, the extra phosphate which appeared in Solution, (Pbicarbonate - Pohloride) = 0.29 If 0.02 corresponded almost exactly to the extra carbonate taken up by the 2 g. of solid, This SUg(CO3solid Clsolid) X 2 = 0.31 f 0.06. gests an equimolar displacement but it is by no means conclusive. From these data, it appears that there is no physicochemical basis for a direct interactlon between sodium and carbonate ions at the surface of hydroxy apatite crystals. Carbonate ions do increase the amount of monovalent ions bound by the solid in a non-specific way, possibly by increasing the charge asymmetry at the crystal :solution interface.16 The close association between sodium and carbonate concentrations in bone seen in vivo must,
-
(14) G. J. Levinskas and W. F. Neuman, THISJOURNAL, 69, 164 (1955). (15) W. F. Neuman and M. Neuman, Chapter IV, “Chemiod
W.
Dynamias of Bone Mineral,” Univ. of Chioago Press, 1958.
NOTES
March, 1958
379
pumping for one hour at 100". Into a cold finger containing 4.6 mmoles of this labeled salt were then distilled 1.62 mmoles of Sic&, using liquid N t summary as a refrigerant, Since HC136, formed by the Both sodium and carbonate ions have previously hydrolysis of Sic14 by any traces of moisture still been shown t o be incorporated in the surface of remaining in the tetramethylammonium chloride, hydroxy apatite crystals by ion exchange processes. would simulate exchange (especially in exchange Because, in bone in vivo, these two ions appear to be runs between labeled SiC14 and unlabeled salt, see related, an attempt was made to test whether any below) the cold finger was momentarily warmed t o specific interaction between sodium and carbonate O o , then cooled with acetone-Dry Ice slush, and ions occurs at the hydroxy apatite surface in vitro. subjected to high vacuum pumping t o remove Carbonate ions did increase incorporation by any HC1 present. The cold finger was warmed to apatite of both sodium and potassium ions but the 0" and aliquots of Sic& removed at known time effects were small and non-specific. The associa- intervals, radioassayed and returned to the reaction between sodium and carbonate ions in bone is, tion cold finger. The counting technique for the therefore, probably the result of physiological assay of volatile labeled materials has been demechanisms, scribed in detail.6 The activity observed in the
therefore, be the result of physiological mechanisms mentioned earlier.
ISOTOPIC EXCHANGE REACTIONS. 11. THE RAPID HALOGEN EXCHANGE BETWEEN Sicla AND (CH&NCl, AND A CONVENIENT METHOD FOR THE PREPARA-
TION OF
c 1 3 6
LABELED CHLOROSILANES
BY ROLFEH. HERBER Department of Chemistry and Chemical Engineering, University of Illinois, Urbana, Ill. Received November I , 1867
In a detailed study1 of the halogen exchange kinetics in the system SiC13nC136-HCln(in which n denotes the natural isotopic composition) it was pointed out that isotopically labeled silicon tetrachloride could be prepared by a homogeneous gas phase exchange with HC136. The reaction halftimes for reactant concentrations of mole per liter were found to be on the order of lo4 seconds at 90". Since, under these conditions the total quantity of SiC14which can be labeled in a given run is small, we have looked for a more suitable exchange path. Schomaker and Stevenson2 have suggested that the difference beiween the sum of the covalent bond radii (2.16 A.) and, the observed interatomic distance in Sic14 (2.02 A.) is due to partial ionic bond character contributions of the type SiC18+C1-. As Lewis and Wilkins3 have shown in the case of NOC1, such ionization should lead to rapid exchange with appropriate ionized solutes and hence we have explored the exchange between (CH&NCl and Sic14 since the former is known to ionize readily in a number of non-aqueous solvent^.^ In a typical exchange run, Cla6-labeled(CH3)4NC1 was prepared by dissolving the reagent grade salt in a minimal volume of water, adding one drop of HC136stock solution5 and evaporating off the water and HC1 under high vacuum conditions. The resulting solid was subjected to high vacuum (1) R. H.Herber, J . Chem. Phys., 27, 653 (1957). (2) V. Schomaker and D. P. Stevenson, J . Am. Chem, Soc., 68, 37 (1941). (3) J. Lewis and R. G. Wilkins, J. Chem. SOC.,56 (1955). (4) A. B. Burg and D. E. McKenzie, J . Am. Chem. Soc., 74, 3143 (1952): V. Gutmann, Z.anow. allgem. Chem.. 266, 331 (1951). (5) Item C1-36-P, Oak Ridge National Laboratory, Isotopes Divi-
sion.
Me,,NCI-S!CI:
AT 30.C.
MINUTES.
Fig. 1.-Activity of silicon tetrachloride.
silicon tetrachloride as a function of time is shown in Fig. 1. In a second experiment, SiC14and labeled (CH3)4NC1maintained for 3 hours at 50°, resulted in an activity of 220 c.p.m./cm. SiCL. We have also noted the reverse exchange in which labeled silicon tetrachloride with an initial specific activity of 200 c.p.m./cm. was contacted with unlabeled (CHs)4NCla at room temperature. The observed c.p.m./cm. were 182 for 5 minutes, 129 after 10 minutes, 77.3 after 20 minutes and 70.7 after 30 minutes. Calculation of the fraction of total exchange, based on the observed "infinite time" activity, appears to indicate that the exchange may proceed via a heterogeneous mechanism. A further experiment in which (CHa)3SiC1nwas substituted for the tetrachloro species indicates that this material, too, exchanges halogens at room temperature with (CH3)4NC136. The use -of high specific activity tetramethylammonium chloride thus provides a conyenient starting point for the preparation of chlorine labeled SiC1, and (CH3)3SiC1and the exchangeis noted to be reasonably rapid at room temperature. Due to the condensed nature of this system, relatively large quantities of labeled product can be prepared in a single exchange run. This research has been supported in part by the U. S. Atomic Energy Commission. The assistance of Mr. W. Cordes in this work is gratefully acknowledged. (6) R . H. Herber Paper #24, 132nd Meeting Am. Chem. SOC., New York City, 1957; R. H. Herber. Rev. 9ci. Inst., in press.
