Effect of Carbonate on Hydroxyapatite Solubility

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Effect of Carbonate on Hydroxyapatite Solubility
  pubs.acs.org/crystalPublished on Web 01/05/2010 r 2010 American Chemical Society DOI: 10.1021/cg901199h 2010, Vol. 10845 – 850 Effect of Carbonate on Hydroxyapatite Solubility Haobo Pan † and Brian W. Darvell* ,‡ † Department of Orthopaedics & Traumatology, The University of Hong Kong, Pokfulam,Hong Kong and  ‡ BioclinicalSciences, Faculty of Dentistry, Kuwait University, Health Sciences Centre,P.O. Box 24923, Safat 13110 KuwaitReceived September 30, 2009; Revised Manuscript Received December 14, 2009 ABSTRACT: Carbonate is present in biological apatites (BAp) (enamel, dentine, bone, and pathological calcifications) bysubstitutionatphosphateandhydroxidesites,tending toincreaseitssolubility incomparisonwithpurehydroxyapatite (HAp).The role of solution carbonate, however, is poorly understood. Using the solid titration method, it was found that the apparentsolubility of HAp increased greatly with an increase in pCO 2 , and was sensitively-dependent at low values. No other phase wasformed at the end point except HAp, and in particular no dicalcium phosphate dihydrate (DCPD) was found. However, nocarbonatewasdetectedintheend-pointsolidexceptforasmallamountatpH ∼ 7.4,pCO 2 =0.01 - 1.0bar.Theimplicationsof these results need further investigation, bearing as they do on a central topic in oral chemistry. Introduction Hydroxyapatite (Ca 5 (PO 4 ) 3 OH; HAp) is generally acceptedtobetheprototypeforthemineralphaseofthecalcifiedtissuesof higher vertebrates, that is, enamel, dentine, and bone. 1 Impurities, especially CO 32 - , are present in these biologicalapatites (BAp) by substitution at either (OH) sites (A-type), orat (PO 4 ) sites (B-type), 2,3 tending to increase its solubility incomparison with pure HAp, and thus believed to enhance thesusceptibility of these tissues to demineralization, for example,dental caries. 4 The role of solution carbonate, however, is stillpoorly understood, and especially in the context of the nuclea-tion and growth of BAp. Although B-type substitution hasbeen considered as most closely resembling that of biologicalapatites, 3,5,6 the substitution extent and the positions of theCO 32 - ions are still controversial; 7 - 9 in particular, highercontent has been reported for immature crystallites such as indentine or young bone (4 - 6 mass%, i.e., equivalent to abouthalf a mol of CO 32 - per formula weight). 10 Previoussolubilitystudies 11 - 21 mayhaveignoredorunder-estimated the effects of carbonate. In fact, as early as 1945,Greenwald 22 had already reported that the solubility of calcium phosphates was increased in the presence of carbo-nate,possiblyduetotheformationofsuchsolutioncomplexesas Ca 2 HPO 4 CO 3 and (Ca 2 PO 4 CO 3 ) - . Later, Ericsson 23 found a much higher solubility for HAp in saliva than inCO 2 -free solution,which again was attributed to the presenceofcarbonicacid.However,evenifunderstood,contaminationby carbonate may account in part for the poor agreement of values reported for the solubility for HAp; 11 - 21 some havetakenstepstoexcludeCO 2 ,othershavenotreporteddoingso.In addition, the conventional “excess solid” method forsolubility determination has recently been demonstrated tobe problematic for calcium phosphates, or at the very leastlacksaccuracy,whilethesolubilityofcarbonate-freeHAphasbeen clearly shown to be substantially lower than widelybelieved. 24 Even so, since carbonate and related species arenormal components of the physiological systems of interest,and saliva in particular; the solubility of HAp in such anenvironment is of fundamental concern if dissolution andprecipitation processes in the context are to be understood.Further detailed study of this system is therefore warrantedand of great importance to oral chemistry and biologicalcalcification in general.Sofar,solidtitration(ST) 25 - 27 hasbeenconfirmedtobeanextraordinarily reliable and reproducible assumption-freemethod for the study of such complex solids and solutions.The use of small solid increments avoids the excessive super-saturationoftheconventionalmethod,inwhichalargeexcessofsolidisaddedinoneportion,withtheseveredisadvantagesthat neither the detailed composition of both solid and solu-tionareknown,northeconstitutionofthesurfaceofthesolidin contact with the solution. Thus, ST can yield a truethermodynamic equilibrium, nucleation permitting, withpre-ciselyknownsolutioncomposition(especiallyCa/Pratio)andavanishinglysmallamountofsolid.Usingthistechnique,twopreviously unreported complexes, 25 CaH 2 PO 4 H 2 CO 3 þ andCaH 2 PO 4 HCO 3 , were postulated from a numerical model toaccount for the increase of calcium-carrying capacity in anartificial saliva equilibrated with 3.5 vol% CO 2 in air(representing a standardized mouth environment). Further-more, a 7-fold increase in solubility was found for HApequilibrated with 3.5 vol% CO 2 in air at pH ∼ 5 in a simple100 mM KClsolution. 26 However,animportant questionliesin whether carbonate is incorporated into the equilibriumsolid, that is, by substitution in the lattice, or only shifts thesolutionequilibriumthroughtheformationofcomplexes.Theaimofthepresentworkwasthereforetocharacterizetheeffectof CO 2 on the mass solubility of HAp and the nature of theequilibrium solid. Materials and Methods Pure HAp solid for use as ST titrant was prepared by a standardprecipitation method through the addition of 0.1 M (NH 4 ) 2 HPO 4 (AnalaR, BDH) dropwise into a 0.167 M Ca(NO 3 ) 2 (AnalaR)solution with the ideal stoichiometric Ca/P ratio of 1.67; solutionpH was controlled by ammonia solution (Aristar, BDH) to around10. Carbon dioxide was excluded by flushing with N 2 . After mixing,the solution was stirred for 2 h, the precipitate collected, washed, *To whom correspondence should be addressed. Fax: þ 965 2498 6698.E-mail: b.w.darvell@hku.hk.  846 Crystal Growth & Design, Vol. 10, No. 2, 2010 Pan and Darvell dried, and finally heated to 800 ° C for 2 h in air. The solid wascharacterized by X-ray diffraction (XRD) (model D/max 2550 V,Rigaku, Tokyo, Japan), using Cu K R (  λ = 1.5406 A ˚) radiationin step-scan mode (2 θ = 0.02 ° per step) (Figure 1) and Fourier-transform infrared spectroscopy (FTIR) (Perkin-Elmer, Wellesley,MA, USA) (Figure 2).Solidtitration,aspreviouslydescribed, 24 - 27 wasusedtodeterminethe mass solubility of HAp in 100 mM aqueous KCl solutionequilibrated with CO 2 in air, partial pressure (pCO 2 ) 0.01 - 1 bar,or nitrogen (for pCO 2 = 0) at37.0 ( 0.1 ° C. The air-CO 2 gases wereobtained premixed and certified (Hong KongOxygen, Hong Kong).Successive portions of pulverized HAp solid were added to thesolution, agitated by a magnetic stirrer. As the end-point wasapproached, these became very small increments (0.5 - 2 mg), as judged appropriate, so as to identify the end-point with precision.The titration process was monitored by the use of a very sensitive,low-angle laser-scattering system to detect the presence of solidparticles. Each addition of solid caused an obvious step-increase inthe scattered laser signal, which then decreased with time due to thedissolution ofthe solid. When a stable signal was obtained at or veryclose tothe originalbaseline, typically foraslong again asittooktheincrementtodissolve(say,2 - 24h),itwastakenasindicatingthattheincrementhadcompletelydissolved,whenthenextadditioncouldbemade. The end-point of the titration was then unambiguouslydetected by the output signal remaining higher than the srcinalbaseline, meaning that no more solid could dissolve, or that a newsolid had precipitated, or both. On achieving stability at this end-point, a further small increment was then made to enable a betterestimate of the actual end-point pH and titrant total by interpola-tion. 27 The pH could then be adjusted by addingdilute aqueous HClsolution (as a noninterfering acid) dropwise such that all solid wasredissolved (normally a deliberately minute amount), as confirmedbythescatteringsignalreturningtothebaseline,againwithsufficientobservationtimetoconfirm,toformthebasisforthenexttitrationtodetermine a new end-point at a lower pH. This reuse procedureenabled an efficient “scan” over a range of pH values. However, toobtainenoughsolidforanalysis(afewmilligrams) atselectedpoints,a few more small increments were added beyond the end-point,ensuringthatequilibrationoccurredforeachone.Such“preparative”precipitates were collected by centrifugation after equilibration for afurther 10 day and characterized by selected-area electron diffrac-tion (SAED) (Technai G2 20 TEM, FEI, Hillsboro, OR, USA)(constitution) since the amount of solid available was inadequate Figure 1. XRD pattern of prepared HAp. The match with thestandard XRD pattern of HAp (JCPDS 72-1243) is good. Figure 2. FTIR spectra of prepared HAp. There is no evidence of carbonate being present. Figure 3. Experimental solubility isotherms of HAp in 100 mMKCl solution equilibrated at pCO 2 0 - 1 bar, at 37.0 ( 0.1 ° C. Figure 4. Extrapolated fitted value at pH 5.5, log(solubility) nearlyproportional to log(pCO 2 ) at pH ∼ 5.5, slope ∼ 0.25.  Article Crystal Growth & Design, Vol. 10, No. 2, 2010 847 for XRD, FTIR (composition), and scanning electron microscopy(SEM) (Leo 1530 FESEM, Oxford, UK) (morphology) at pH ∼ 3.8and 4.5. ST could be used effectively only up to pH ∼ 6, when thesolubility became too low for the resolution of the system. However,precipitates were prepared at pH ∼ 7.4 by a similar method of suc-cessive addition of very small increments, in order toobtain materialrelevant to the normal human internal physiological environment.Thiswaspredicatedontheobservationthatinallpriorwork,srcinalsolidtitrantwasneverobserved(SEM)inequilibratedsolid(asmightordinarilybeexpectedtobefoundaccordingtotheconventionalideaof solubility), it being presumed that the titrant was sufficientlydamagedbypulverization, atleast,astorenderitunstableincontactwith the environmental solution. Results The titrant showed no detectable presence of other phases(Figure 1), and in particular there was no evidence forcarbonate (Figure 2). The mass solubility of HAp wasgreatly increased with the increase in pCO 2 at high pH( c . 5.4) (Figure 3), with convergence toward a common value(within the resolution of the system) near pH 3.5. Figure 4shows that the log(solubility) is nearly proportional to log-(pCO 2 ) at pH ∼ 5.5, using fitted polynomial lines to inter-polate or extrapolate the values (Figure 3). No phase otherthan HAp was detectable in the precipitates at either pH ∼ 3.8, 4.5 or 7.4, judging by the calculated d  -spacings, whichmatch well the standard values for HAp (JCPDS 72-1243)(Figure 5; Table 1). The FTIR results show that no C - Ostretching bands, assigned to the CO 32 - group, at about1490, 1415, and 872 cm - 1 28 were detectable; that is, therewas no evidence of CO 3 substitution, at either pH 3.8 or4.5. However, CO 3 was weakly detectable, with a broad bandfrom 1410 to 1450 cm - 1 , in the solid prepared by equilibra-tion at pH 7.4, where the material appeared slightly amor-phous(Figure6)andthecalculated d  -spacingsdeviateslightlyfrom the standard values (Table 1), possibly due to theincorporationofcarbonate,althoughthecontentwasdifficultto determine. Discussion The absence of detectable CO 3 substitution in the equili-bratedsolidatpH3.8and4.5isaclearindicationthat,despitethe increased mass solubility, at least over the range up to Figure 5. SAED rings for precipitates at pCO 2 (a) 0.01 bar (b) 1 bar. Ring numbers as in Table 1. Table 1. The Calculated d  -Spacings for the Selected-Area ElectronDiffraction ( SAED ) Rings of Figure 5 at pCO 2 ( a ) 0.01 bar ( b ) 1 bar a (a)calculated d  -spacings/A ˚HApring pH 3.84 pH 4.51 pH 7.38 d  /A ˚ hkl  1 3.4563 3.4517 3.4716 3.4405 0022 3.1831 3.1803 3.1903 3.1707 1023 2.8217 2.8198 2.8296 2.8168 2114 2.3018 2.3011 2.3077 2.2978 2125 1.9489 1.9467 1.9513 1.9450 2226 1.9003 1.8997 1.9029 1.8921 1327 1.8477 1.8473 1.8498 1.8411 213(b)calculated d  -spacings/A ˚HApring pH 3.87 pH 4.68 pH 7.41 d  /A ˚ hkl  1 3.4531 3.4507 3.4703 3.4405 0022 3.1827 3.1798 3.1912 3.1707 1023 2.8203 2.8201 2.8287 2.8168 2114 2.2979 2.3007 2.3071 2.2978 2125 1.9473 1.9486 1.9518 1.9450 2226 1.8960 1.8958 1.9012 1.8921 1327 1.8456 1.8483 1.8411 2138 1.7233 1.7202 004 a The values for HAp (JCPDS 72-1243) are shown for comparison.The radiiwere calculatedin imageanalysissoftware (LeicaQWin,LeicaMicrosystems Imaging Solutions Ltd., Cambridge, UK).  848 Crystal Growth & Design, Vol. 10, No. 2, 2010 Pan and Darvell ∼ pH 5.5, HAp is the equilibrium solid despite the presence of CO 2 in solution. This has three important connotations.First, the exclusion of CO 2 is critical for the determinationof the mass solubility of HAp in particular, and calciumphosphates in general, whatever approach is employed. Nodata for any such system where CO 2 has not been explicitlyand rigorously excluded can be considered reliable, as even atthe pCO 2 of air an appreciable increase in mass solubility isexpected(Figure4),andlackofcontrolmustleadtoincreasedscatter as well. As an indication of this, attempts to obtaingood results for the natural air-equilibrated system (pCO 2 ∼ 0.035) proved frustrating due to the slowness of that equili-bration, yet while noisy these data (not shown) were essen-tially consistent with the behavior now presented. There is noreasontoexpectaqualititativechangeundertheseconditions.Second, the ordinary thermodynamic expectation of themost stable (i.e., least soluble) solid has been met. As indi-cated, carbonated apatite is believed to be more soluble thanHAp;thus,itshouldnotformeveninthepresenceofsolutioncarbonatewhenthesolubilityproductofthemorestableHApis attained. To do so would violate a strong rule unless therewereanucleationorkineticlimitation.Thus,evenifthetitrantwere to be contaminated by carbonate, even undetectably, itcouldhavenoeffectorbearingonanysolubilityresultsinthispH range when continuous flushing with an equilibratinggas was used (providing, of course, that sufficient time wasallowed).Third, the evident increase in the mass solubility of HAp(which may be viewed as the calcium-carrying capacity of thesolution) must unambiguously be entirely due to the forma-tion of solution complexes involving carbonate - there is noother possibility. What those complexesarecanbe left for themoment,butitfollowsimmediatelythatunlesstheyaretakenintoaccount,fullyand accurately, noattemptat calculatingasolubility product for any calcium phosphate, and HAp inparticular, will be meaningful if CO 2 is present, for example,from exposure to the air.Greenwald 22 was working in the pH range 6.9 - 9.0, but asobserved by Leung & Darvell, 25 the species then proposed -CaHPO 4 CO 32 - and CaPO 4 CO 33 - - only differ in terms of protonation, so that the extra possibilities of CaH 2 PO 4 -H 2 CO 3 þ , CaH 2 PO 4 HCO 3 , and CaHPO 4 HCO 3 - form partofacontinuousseries(althoughitisnotpossibletodistinguishbetween such forms as CaHPO 4 H 2 CO 3 þ and CaH 2 PO 4 -HCO 3 , the effect is the same as there will be an equilibriumbetween the two). In other words, a series of Ca-CO 3 -PO 4 species are expected to exist, varying only by the degree of protonation, from at least 4 to 0, depending on pH. This is anaturalsequence that hasbeensuggested toaccount for,interalia, the apparent supersaturation of saliva with respect toHApratherthanjustthesimplisticbufferingcapacityconcept(loss of CO 2 causes precipitation) usually invoked. 29 Thepresent results are entirely consistent with the existence andoperation of such a system. The fact that the apparentsolubility is independent of pCO 2 for pH < ∼ 3.5 (Figure 3),is taken to indicate near-complete complexation of Ca withcarbonate and phosphate at low pH, in other words CO 2 seems to be absorbed to titrate fully the Ca present, whichin turn implies high stability for the relevant complex(es).Numerical modeling is required to resolve this fully.Again,ithasbeenfound,consistentwithpreviousSTwork,that HAp is the most stable phase at pH < 4.2, despite otherclaims and reports that dicalcium phosphate dihydrate(DCPD) is and should be the equilbrium phase; 30 indeed, ithas been clearly shown to be metastable, 31 - 33 dissolving infavor of HAp on seedingwithHAp. 33 Thus,the conventionalexcess-solid method is simply not suitable for calcium phos-phate studies because it would appear that the great super-saturation it engenders leads, in an Ostwald succession sense,to the metastable formation of DCPD, at least in that pHregion. This conflict between the thermodynamic expectationand the kinetic reality has probably been the cause of muchmisunderstanding of this system. However, more critical isthat false solution equilibria may have been derived on thebasis of such data, and used erroneously to explain manyphenomena both in vitro and in vivo over a long period.Thepresenceofcarbonateintheprecipitateswas,however,easilydetectablewhenthepHwasincreasedtoaphysiological-like value, but this presence appeared to be independentof pCO 2 (Figure 6), which is at odds with a conventionalthermodynamicallybased,solubilityproduct-determinedout-come.Simply,ifacarbonatedapatiteexistsasadistinctphase,it should form in an all-or-nothing fashion when conditionsare such as to make it the most stable phase, or (again in the Figure 6. FTIR spectra for precipitated solids in titrations at pCO 2 (a) 0.01 bar (b) 1 bar.  Article Crystal Growth & Design, Vol. 10, No. 2, 2010 849 Ostwald succession sense) asa metastable intermediate whichwould transform on equilibration to the next most stablephase (presumably HAp itself). Even if, as seems morelikely, carbonated apatite is a member of a solid - solutioncontinuum (as is the case for fluoride, for example, 34 andmanyother“contaminants”),apCO 2 -dependenteffectwouldstill be expected. In nature, no pure form of apatite can befound: substitutions for Ca 2 þ (Sr 2 þ , Ba 2 þ , Mg 2 þ , K þ , Na 2 þ ,Fe 3 þ ), PO 43 - (AsO 43 - , CO 32 - , VO 43 - ), and OH - (F - ,CO 32 - , Cl - , Br - ) all occur. 35 Indeed, “biological apatite” iscarbonate-apatite with traces of sodium, potassium, andmagnesium.While the partial replacement of PO 43 - or OH - (or both)byCO 32 - musthaveoccurred,itisnotpossiblewiththesedatato distinguish between A- and B-type substitution, althoughtheOH - stretchingbandappearstohavebecomeweaker.Thelarger particles seen at pH 7.4 (Figure 7) may be attributed to“Ostwaldripening”,butjudgingfromthemorediffuseSAEDpattern(Figure5)andtheFTIRband-blurring(Figure6),thecrystallinity was poorer. Consequently, it seems clear thatwhile the formation of carbonate-apatite is not thermo-dynamically favored, it may be that it nucleates and growseasily (which suggests disorder) at pH ∼ 7, but the equilibra-tion time allowed now was insufficient.Inanycase,failuretounderstandorsimplyoverlookingtheeffect of carbonate must lead to a false result, just as wouldcontaminantssuchasfluoride. 34 However,itmustbesaidthatunder physiological conditions, and especially in the complexenvironments in which biominerals are formed, the conflictbetween thermodynamics and kinetics is likely to mean thatsimple equilibria will not be a sufficient explanation. Solu-tion speciation, and especially for CO 32 - -containing species,requires more detailed study and numerical modeling.Insummary, the dramatic effect of solution carbonate maybe attributed to the formation of previously postulated com-plexes such as CaH 2 PO 4 H 2 CO 3 þ and CaH 2 PO 4 HCO 3 shift-ing the equilibrium in favor of greater dissolution. It followsthat previously reported solubility data may need re-evalua-tioninthelightoftheeffectof“contamination”levelsofCO 2 .Nevertheless, the stable phase is HAp, at least at low pH.While the equilibrium carbonate content at physiological pHappears to be very limited, this presence may be a kineticrather than a thermodynamic phenomenon. Acknowledgment. We are grateful for the assistance of Ms. Xiaoli Zhao of the Department of Orthopaedics &Traumatology, The University of Hong Kong, in obtainingthe SAED data and its analysis. References (1) LeGeros, R. Z. Calcium Phosphates in Oral Biology and Medicine ;Karger: Basel, 1991. (2) Leventouri, Th.; Chakoumakos, B. C.; Moghaddam, H. Y.;Perdikatsis, V. J. Mater. Res. 2000 , 15 , 511  –  517. (3) Wilson, R. M.; Elliott, J. C.; Dowker, S. E. P.; Smith, R. I. Biomaterials 2004 , 25 , 2205  –  2213. (4) Cury, J. A.; Francisco, S. B.; Sim ~ oes, G. S.; Del Bel Cury, A. A.;Tabchoury, C. P. M. Caries Res. 2003 , 37  , 194  –  199. (5) Elliott, J. C. Structure and Chemistry of the Apatites and OtherCalcium Orthophosphates ; Elsevier: Amsterdam, 1994. (6) LeGeros, R. Z. Crystallography studies of the carbonate substitu-tion in the apatite structure, PhD Thesis, New York University,New York, 1967 . (7) Leventouri, Th.; Chakoumakos, B. C.; Moghaddam, H. Y.;Perdikatsis, V. J. Mater. Res. 2000 , 15 , 511  –  517. (8) Leventouri,Th.; Moghaddam, H. Y.;Papanearchou, N.; Bunaciu,C. E.; Levinson, R.; Martinez, O. Mater. Res. Soc. Symp. Proc. 2000 , 599 , 79  –  84. (9) Ivanova, T. I.; Frank-Kamenetskaya, O. V.; Kol’tsov, A. B.;Ugolkov, V. L. J. Solid State Chem. 2001 , 160 , 340  –  349. Figure 7. Precipitate morphology by SEM at pCO 2 (a) 0.01 bar (b) 1 bar.
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