Crystallization of 4′-Hydroxyacetophenone from Water: Control of

May 17, 2012 - The crystallization precedence of the various phases, their approximate lifetimes, and transformation sequences, while in contact or wh...
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Crystallization of 4′-Hydroxyacetophenone from Water: Control of Polymorphism via Phase Diagram Studies Carlos E. S. Bernardes†,‡ and Manuel E. Minas da Piedade*,† †

Departamento de Química e Bioquímica e Centro de Química e Bioquímica, Faculdade de Ciências, Universidade de Lisboa, 1649-016 Lisboa, Portugal ‡ Centro de Química Estrutural, Complexo Interdisciplinar, Instituto Superior Técnico da Universidade Técnica de Lisboa, 1049-001 Lisboa, Portugal S Supporting Information *

ABSTRACT: The preparation of polymorphs and solvates and the characterization of their stability domains have received considerable attention in recent years, due to the importance of these studies for fundamental research and for the production of new materials for task-specific applications. In this work, the selective and reproducible crystallization of different solid forms of 4′-hydroxyacetophenone (HAP) from water was investigated, through the determination of a temperature−concentration (T− cHAP) phase diagram. This determination was mainly based on gravimetric solubility measurements, slurry tests, and metastable zone width (MZW) studies with thermometric and turbidity detection. The experimental conditions for the formation of five different HAP phases by cooling crystallization could be established: the previously characterized anhydrous forms I and II and the hydrate HAP·1.5H2O (H1), and two new hydrates, one of stoichiometry HAP·3H2O (H2) and another (H3) which proved too unstable for a stoichiometry determination. The crystallization precedence of the various phases, their approximate lifetimes, and transformation sequences could also be elucidated. It was finally found that for a specific T−cHAP domain the crystallization of HAP solid phases was mediated by a colloidal dispersion. Preliminary dynamic light scattering experiments indicated that this dispersion consisted of particles with diameters in the range of 100−800 nm.

1. INTRODUCTION The study of polymorph and solvate occurrence in molecular organic solids has received considerable attention in recent years due to the opportunities offered in terms of fundamental research and production of new materials with commercial value.1−4 Because dissimilar polymorphs have different packing architectures, they can also exhibit significant differences in properties, such as the fusion temperature, compressibility, or the solubility and dissolution rate in a given media. The control of polymorph formation is, therefore, of considerable interest to the fine chemicals industry (e.g., pharmaceuticals) since it provides a means to tune the properties of a product in view of an application, without changing the molecule involved. This concept can also be extended to solvates when the presence of solvent in the crystal lattice can be tolerated.1−4 © 2012 American Chemical Society

Two major issues normally arise after polymorphs or solvates have been isolated and structurally characterized: the definition of their thermodynamic stability domains and kinetic barriers for phase transitions or thermal decomposition, and the development of methodologies for their selective and reproducible preparation. Indeed two or more polymorphs or solvates can often be prepared and stored at ambient temperature and pressure (∼298 K and ∼1 bar), but in the absence of kinetic barriers any metastable form will transform over time into the thermodynamically stable modification under those conditions.1−4 It is also frequent that new polymorphs or solvates are obtained by accident, or in the course of crystallization screening tests without tight control of Received: January 30, 2012 Revised: May 16, 2012 Published: May 17, 2012 2932

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5−35°, with 2θ step sizes of 0.015° or 0.017° and scan step times of 1.0 or 20 s for the first and second instruments, respectively. The samples were mounted on aluminum sample holders. The program Checkcell10 was used for the indexation of the powder patterns and the diffractogram simulations were performed by using the Mercury 2.311 software package and the previously reported single crystal X-ray diffraction data for the anhydrous HAP phases.6,7 Optical microscopy images were obtained with an Olympus SZX10 stereoscopic microscope and the CellD 2.6 software. The 4′-hydroxyacetophenone (Aldrich, mass percentage 98%) used as starting material for all the solubility and crystallization experiments was purified by sublimation at 368 K and 13 Pa. This procedure was previously shown by gas chromatography−mass spectrometry (GCMS) analysis to typically lead to samples with a mass percentage purity >99.9%.6 The powder pattern of the purified material was indexed as monoclinic, space group P21/c, with a = 7.675(4) Å, b = 8.354(18) Å, c = 11.210(6) Å, β = 95.09(7)°. These results are consistent with those obtained for form I HAP, by single crystal X-ray diffraction measurements carried out at 298 K: a = 7.7200(15) Å, b = 8.3600(17) Å, c = 11.280(2) Å, and β = 95.02(3)°.6 Distilled and deionized water from a Millipore system (conductivity ≤0.1 μS) was used in the preparation of all solutions. In addition to water, ethanol (Panreac; mass percentage 99.9%) was also used in the preliminary analysis of the relative thermodynamic stability of HAP forms I and II based on slurry studies. 2.2. Solubility Measurements by the Gravimetric Method. The determination of the solubility of HAP in water in the range 275− 323 K by the gravimetric method essentially followed a previously described procedure.12,13 A suspension of HAP in ∼200 cm3 of water was magnetically stirred, under nitrogen atmosphere, inside a 250 cm3 jacketed Schlenk glass cell. The temperature of the mixture was controlled to ±0.05 K by circulating water from a JULABO F25-EC thermostatic unit through the outer jacket of the cell and monitored with an accuracy of ±0.01 K by using a Pt100 sensor connected in a four wire configuration to an Agilent 34970A multimeter. The sensor had been previously calibrated against a reference platinum resistance thermometer, calibrated at an accredited facility in accordance to the International Temperature Scale ITS-90. The mixture was initially maintained at 283.25 ± 0.05 K for at least 24 h. Stirring was then stopped and samples of the saturated solution (∼2 cm3) were collected in triplicate by using a syringe adapted to a Schleicher & Schuell BA 65 membrane filter (45 μm, i.d. 25 mm) and a Hamilton 7748-06 stainless steel needle. To avoid precipitation the syringe, filter, and needle were, prior to use, kept in an oven, whose temperature was adjusted to ∼10 K above the temperature of the mixture. The sample aliquots were transferred to preweighed glass vials of 10 cm3 volume, which were weighed a second time when loaded with the solution and a third time after the solution had been taken to dryness. These series of weighings, performed with a precision of ±0.01 mg on a Mettler Toledo XS205 balance, led to the concentration of HAP in each solution aliquot and the mean of the three results was taken as the concentration of HAP in the saturated solution. The solubility determinations were performed at ascending temperatures, separated by ∼5−10 K, up to 322.86 K. If necessary, to ensure saturation, further solid was added to the system as the temperature increased. The procedure was repeated on descending the temperature to 275.23 K. In the course of these experiments the appearance of a “milky” colloidal dispersion was noted on heating the HAP + H2O mixture slightly above the upper temperature limit of ∼323 K. The formation of this dispersion as a precursor of the crystallization of different HAP solid phases from aqueous solution will be discussed below. 2.3. Cooling Crystallization Studies. MZWs for cooling crystallization of HAP from water in the HAP concentration range 16 g·kg−1 < cHAP < 60 g·kg−1 (cHAP in grams of anhydrous HAP per 1 kg of water) were measured by using the in-house designed CB1 reactor and the operating procedure previously reported.12 The core of this system essentially consisted of a glass vessel closed by a steel lid. The lid supported a turbidity immersion probe (Avantes FCR71R200-2-45-ME), a thermometer (a PT100 sensor connected in a

the experimental conditions. In these cases the preparations may be difficult to replicate, and this has motivated numerous efforts to develop systematic methodologies for the selective and reproducible production of polymorphs and solvates.1−4 Crystallization from solution is by far the most widely used method for the preparation of a product corresponding to a specified crystalline form with consistent and well-characterized physical properties.1−5 In a previous work this method (with water or ethanol being the solvent) was used to obtain two anhydrous modifications (form I, monoclinic, P21/c, Z′ = 1; form II, orthorhombic, P212121, Z′ = 2)6 and one hydrate form (H1; triclinic, P1̅ Z′ = 1)7 of 4′-hydroxyacetophenone (HAP), which were then characterized in terms of structure and thermodynamic stability domains. The thermodynamic study of the anhydrous polymorphs indicated that form II was the most stable at 298.15 K. It also revealed that the two phases were enantiotropicaly related: differential scanning calorimetry (DSC) experiments consistently evidenced an endothermic form II → form I transition at 351.2 ± 2.7 K, which was followed by fusion of form I at 381.9 ± 0.1 K.6 The solid−solid phase transition could only be observed on heating the sample. It was, therefore, impossible to establish if the transition temperature obtained by DSC corresponded to the true equilibrium value and this problem was addressed in the present work. The hydrate of stoichiometry HAP·1.5H2O had the water molecules lined up in chains that resided in lattice channels (channel hydrate). Dehydration readily occurred when the solid was removed from contact with the mother liquor and led to the anhydrous form I. A thermodynamic analysis of the dehydration process indicated that, at 298.15 K, the loss of water became favorable (ΔrGm < 0) for a relative humidity Φ < 66%, with the corresponding temperature increasing to 329 K for the saturation limit Φ ∼ 100%.7 Although the reported preparation recipes consistently yielded the desired phases, no study of the underlying crystallization processes was carried out. If, however, tight control over selective crystallization of the different HAP forms is to be achieved, then a reasonable understanding of their nucleation and crystal growth is necessary.5,8 A key element toward this goal is a temperature vs concentration (T−c) phase diagram illustrating the saturation and supersaturation limits (metastable zone width) of HAP in the solution, and the stability domains of the different crystallized phases.5,8,9 In this work the T−c phase diagram of the HAP + H2O system was investigated based on solubility and metastable zone width (MZW) determinations and on the characterization of the crystalline phases obtained for different T−c ranges. The main objective was to achieve a better control over the selective production of 4′-hydroxyacetophenone solid forms by cooling crystallization and to gain some insight into their interconversion relationships while in contact with the solvent.

2. MATERIALS AND METHODS 2.1. General. X-ray powder diffraction (XRPD) analyses were carried out at 296 ± 3 K in two apparatus: a Philips PW1730 equipped with a vertical goniometer (PW1820), a proportional xenon detector (PW1711), and a graphite monocromator (PW1752) or in a Philips Analytical X’Pert PRO, set with a vertical PW 3050/60 goniometer, and a X’Celerator detector. Automatic data acquisition was performed by means of the APD Philips v.35B and the X’Pert Data Collector v2.0b software packages, respectively. Both instruments operated in the θ−2θ mode, using Cu Kα radiation sources, with the tube amperage set to 30 mA and the tube voltage to 40 kV. The diffractograms were recorded by continuous scanning in the 2θ range 2933

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was controlled by a JULABO F33-ME thermostatic unit that allowed the programming of constant rate heating/cooling ramps. 2.4. Thermogravimetry (TG). The TG experiments on the unstable H2 hydrate mentioned below were carried out on a PerkinElmer TGA 7 apparatus. The balance chamber was kept under a positive flow of nitrogen (Air Liquide N45) of 38 cm3·min−1. The sample purge gas was helium (Air Liquide N55) at a flow of 22.5 cm3·min−1. The mass scale of the instrument was calibrated with a standard 100 mg weight and the temperature calibration at 5 K·min−1 was based on the measurement of the Curie points (TC) of alumel alloy (Perkin-Elmer, TC = 427.35K) and nickel (Perkin-Elmer, mass fraction 0.9999, TC = 628.45 K) standard reference materials. The hydrate phase was separated from the mother liquor by means of vacuum filtration using a sintered glass disk filter funnel. A sample of the obtained material with an initial mass of 8−25 mg was transferred to the open platinum crucible of the TG apparatus and studied in the temperature range 298−373 K at a scan rate of 5 K·min−1. Because of the unstable nature of the hydrate its isolation and transfer to the TG balance had to be performed as quickly as possible. In general, these operations were completed in ∼1 min.

four wire configuration to an Agilent 34970A multimeter), a stirring system, and a needle inlet for automatic addition of solvent from a buret (Crison Buret 1S; precision better than ±1 μL). Before entering the reactor, the solvent traveled through 2.5 m of Teflon tube immersed in the thermostatic bath. This ensured that water was added at a known temperature and allowed the calculation of the corresponding mass, m/g, through the equation:12

m = V (− 1.44510 × 10−10T 4 + 2.02215 × 10−7T 3 − 1.08687 × 10−4T 2 + 2.59474 × 10−2T − 1.29515) (1) 3

where V/cm is the volume dispensed by the buret and T/K is the temperature of the thermostatic fluid at the time of the water addition. The scattered intensity was simultaneously analyzed at 550, 600, and 650 nm by means of an Avantes AvaSpect 2048 spectrophotometer linked through fiber optics to the turbidity sensor. A stirring rate of 420 rpm was used. The temperatures corresponding, on cooling, to the onsets of crystallization (Tc) or formation of the “milky” colloidal dispersion mentioned above (Td), and, on heating, to the complete dissolution of the solid (saturation temperature, Ts), were determined from the corresponding singularities observed in the turbidity-time and/or temperature−time profiles. The glass vessel was initially loaded with ∼2.8 g of HAP and ∼47.0 g of H2O (cHAP ∼ 60 g·kg−1). Both the solvent and the solute were weighed inside the vessel with a precision of ±0.1 mg, by using a Mettler Toledo XS205 balance. The steel lid was adapted to the glass vessel and the assembled reactor was transferred to a thermostatic bath (JULABO F33-ME) whose temperature could be varied at a programmed rate. Stirring was started and the mixture was subjected to a series of heating and cooling cycles at rates β = 12 K·h−1, 18 K·h−1, and 24 K·h−1. The programmed experimental conditions had to be slightly different when the runs crossed temperature−concentration ranges where the colloidal dispersion was present and when they only passed through zones where the dispersion did not occur. If the dispersion was formed on heating the mixture, no singularities in the turbidity and temperature time profiles could be detected at the saturation temperature, Ts. The turbidity probe was, however, able to detect its formation on cooling at Td. Furthermore, to ensure that a homogeneous solution had been generated before the start of a cooling ramp, all heating steps were ended at 342 K. This value was selected based on preliminary tests carried out with visual detection using a jacketed Schlenk glass cell (see below). Thus, for runs where the colloidal dispersion occurred, the mixture was initially heated at 24 K·h−1 to 342 K, kept at this temperature for 20 min, and then subjected to the following thermal analysis program: (i) cooling ramp at 24 K·h−1 to 5 K below the onset of crystallization, Tc, which was signaled by slope changes in both the turbidity−time and temperature−time profiles; (ii) 20 min isothermal step; (iii) heating ramp at 24 K·h−1 to 342 K; (iv) 20 min isothermal step. After performing identical cycles at 12 K·h−1 and 18 K·h−1, a volume of 2.0−2.5 cm3 of water was automatically added to the mixture from the buret and the entire sequence was repeated for the new concentration of the mixture. For temperature−concentration domains where the dispersion did not form, the onset of complete dissolution could be automatically detected by the turbidity sensor and the above step (iii) was replaced by a heating ramp at a selected rate to 5 K above Ts. The dilutions were performed until the HAP concentrations reached cHAP ∼ 16 g·kg−1. To elucidate the nature of the crystallization events detected with the CB1 reactor, the cooling crystallization of mixtures with some specific HAP concentrations was also studied, under magnetic stirring, in a 100 cm3 jacketed glass vessel. This set up allowed sampling the crystalline phases immediately after precipitation and their subsequent characterization in terms of crystal form and morphology, and stoichiometry in the case of hydrates. The temperature inside the vessel was monitored with an accuracy of ±0.5 K with a mercury thermometer and could be varied at 18 K·h−1 by circulating thermostatted water through the outer jacket. The water temperature

3. RESULTS AND DISCUSSION All HAP concentrations, cHAP, refer to grams of anhydrous HAP per 1 kg of water. 3.1. Equilibrium Solubility Studies. Solubility Curves. The solubility of HAP in water was determined by the gravimetric method in the range 275−323 K. The obtained results using form I as starting material are illustrated in Figure 1 (see Supporting Information for source data). The

Figure 1. Solubility of 4′-hydoxyacetophenone in water obtained on ascending (○) and descending (●) the temperature. The concentration of HAP, cHAP, refers to grams of anhydrous solute per 1 kg of water.

observation of a kink at ∼302 K in the solubility curve of Figure 1 suggested that different materials might be in contact with the solution below and above this temperature. This was confirmed in a series of tests described below, which led to the conclusion that for T < 302 K form I HAP consistently transformed into a hydrate phase (here dubbed H2) during the equilibration period. Above 302 K the hydrate H2 decomposed to yield form I HAP. The formation of H2 for T < 302 K and its transformation into form I for T > 302 K was also found to take place if form II HAP was initially used. The solubility of HAP in water from 275.2 to 302 K (Figure 1, H2 domain) and from 302 to 322.9 K (Figure 1, form I domain) is given by the equations: c HAP = 1322.37 − 9.48119T + 0.01704T 2 (275 − 302 K) 2934

(2)

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Figure 2. Images of the HAP crystalline phases at different times, recorded at ∼296 K: (a) original form I HAP before contact with water; a sample of the solid removed from the solution after (b) 30 min and (c) 150 min equilibration time at 296 K, respectively; (d) the product in image (c) after being kept in air for ∼20 min; and (e) the solid removed from the solution after a 48 h equilibration time at 302.6 K.

K, respectively; (iii) the product in Figure 2c after being kept in air for ∼20 min (Figure 2d); and (iv) the solid in contact with the solution after a 48 h equilibration time at 302.6 K (Figure 2e). The fact that the crystals in Figure 2a,b were stable in air and those in Figure 2c turned opaque in ∼20 min (Figure 2d) suggested that the solid material in contact with the solution evolved into a hydrate phase, which becomes unstable if removed from the solution. This was confirmed by TG experiments carried out on samples collected from the slurry after ∼24 h contact. The results indicated that incipient water loss already started while the material was being transferred to the TG apparatus (no steady initial baseline could be observed) and that the compound stoichiometry was approximately HAP·3H2O (see Supporting Information for details). Although crystals of this hydrate (labeled above as H2) suitable for a structural determination by single crystal X-ray diffraction could not be grown, further evidence that it was different from the previously reported HAP·1.5H2O phase (H1)7 was provided by a comparison of their XRPD patterns (Figure 3). The XRPD

c HAP = 4316.05 − 28.7160T + 0.04791T 2 (302 − 323 K)

(3)

where cHAP/g·kg−1 is the concentration of HAP in solution and T/K is the absolute temperature. Equations 2 and 3 were based on a least-squares fitting to the experimental cHAP against T data (Figure 1). Both fittings gave a determination coefficient R2 = 0.99 for 95% probability. The uncertainties assigned to the calculated cHAP values are σ = ± 0.44 g·kg−1 and σ = ± 0.88 g·kg−1 for eqs 2 and 3, respectively. The σ values represent standard deviations and were derived from the differences between the corresponding cHAP experimental and calculated results. Characterization and Stability of Solid Phases. Figure 2 shows a sequence of images recorded at ∼296 K for (i) the original form I HAP before contact with water (Figure 2a); (ii) a sample of the solid removed from the solution after 30 min (Figure 2b) and 150 min (Figure 2c) equilibration time at 296 2935

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Figure 3. Comparison of the X-ray powder diffractograms of the previously reported HAP·1.5H2O hydrate (H1)7 and the HAP·3H2O hydrate (H2) observed in this work. Both patterns refer to 296 ± 3 K.

Figure 5. Comparison of the experimental X-ray powder diffractograms of the solid removed from the solution after a 48 h equilibration time at 302.6 K (Figure 2a) and that of form I HAP simulated from published6 single crystal X-ray diffraction results. Both patterns refer to 296 ± 3 K.

experiments further showed that the dehydration of H2 crystals in air, at ambient temperature (∼296 K), leads to the anhydrous form II HAP and not to form I HAP as when heated above ∼302 K in contact with the solution. This is evidenced in Figure 4, where the powder pattern of dehydrated

The hydrate H2 slowly evolved into a new phase if the solid suspension was kept under magnetic stirring, at ∼293 K. This is shown in Figure 6 where the X-ray powder diffractograms of

Figure 6. X-ray powder diffraction patterns recorded over a six month period for the solid in contact with the HAP aqueous solution at ∼293 K. The data for 180 days are compatible with that for the original H1 hydrate phase presented in Figure 3. The XRPD patterns refer to 296 ± 3 K.

Figure 4. Comparison of the experimental X-ray powder diffractograms of dehydrated H2 and of form II HAP simulated from published6 single crystal X-ray diffraction results. Both patterns refer to 296 ± 3 K.

H2 is compared with that of form II HAP simulated from published single crystal X-ray diffraction results obtained at 298 K.6 It was further supported by the corresponding indexation (orthorhombic, space group P212121, with a = 6.1296(6) Å, b = 9.5359(6) Å, c = 24.3289(22) Å; see Supporting Information for details) which was in good agreement with the characterization of form II HAP by single crystal X-ray diffraction (orthorhombic, space group P212121, a = 6.1097(11) Å, b = 9.5293(14) Å, c = 24.313(4) Å).6 In contrast, the dehydration of H1 in air was previously observed to yield form I HAP.7 The solid phase resulting from decomposition of H2 while in contact with the solution above ∼302 K (Figure 2e) was shown to be anhydrous by thermogravimetry (mass loss of water smaller than 0.03% in the range 298−363 K) and subsequently assigned to form I HAP based on XRPD analysis (Figure 5). Indeed the corresponding powder pattern could be indexed as monoclinic, space group P21/c, with a = 7.6951(9) Å, b = 8.3141(11) Å, c = 11.2577(16) Å, β = 95.06(1)° in excellent agreement with the corresponding data for form I HAP previously obtained from single crystal X-ray diffraction experiments, namely, monoclinic, space group P21/c, a = 7.7200(15) Å, b = 8.3600(17) Å, c = 11.280(2) Å, and β = 95.02(3)°.6

samples collected from the mixture during six months are compared. The results suggest that the transformation leads to the H1 hydrate, whose original powder pattern is shown in Figure 3. The observation of a very slow H2 → H1 transition highlights the fact that the solid−solution equilibrium may be difficult to establish, within the equilibration periods typical of solubility studies (a few hours to a few days). It also indicates that, in spite of its usefulness from a practical application point of view, Figure 1 does not correspond to a true equilibrium phase diagram. The time evolution of the HAP + H2O system discussed in this section, above and below 302 K, is summarized in Scheme 1. Form II → Form I Transition Temperature. The finding that from ∼302 K (the upper temperature limit of the H2 domain in Figure 1) to 323 K (the upper temperature limit of solubility measurements), form I and not form II was in equilibrium with the solution, is in apparent contradiction with the stability domains of the two anhydrous polymorphs of HAP suggested by their previous calorimetric study.6 Indeed, solution calorimetry and drop-sublimation Calvet microcalorimetry results led to the conclusion that form II was more stable than form I at 298.15 K and DSC experiments complemented 2936

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Scheme 1

by XRPD analysis evidenced the presence of an endothermic and irreversible form II → form I phase transition at 351.2 ± 2.7 K, indicating that form II might be more stable than form I up to this temperature. To further investigate this issue, a preliminary slurry transformation test was carried out where crystals of HAP forms I and II were independently kept in contact with ethanol for over one week, at 293 and 308 K, respectively. The replacement of water by ethanol was necessary due to the formation of the H2 hydrate below 302 K in aqueous media. The obtained results support the observations made for the HAP + water system: form II was found to convert into form I after being kept under magnetic stirring for ∼24 h at 308 K and, conversely, the form I → form II transition was completed in ∼1 week at 293 K. The slurry tests therefore indicate that, as previously suggested by the solution calorimetry and drop-sublimation Calvet microcalorimetry experiments, form II is more stable than form I at 298.15 K. The equilibrium form II → form I transition seems, however, to occur at a much lower temperature than judged from the DSC results (351.2 ± 2.7 K). This is consistent with Burger’s enthalpy of transition rule, which implies that if an endothermic solid−solid phase transition is observed in a DSC trace, then the corresponding equilibrium temperature must be at, or below, the temperature of the experimentally detected peak.3,14,15 The accurate determination of the equilibrium temperature of the form II → form I transition is currently under way and will be the topic of a subsequent publication. 3.2. Cooling Crystallization. Phase Diagram. The phase diagram summarizing the results of the MZW determination is illustrated in Figure 7 (see Supporting Information for source data). Also included in Figure 7 are the solubility data obtained by the gravimetric method. On average, these agree with the corresponding values from the MZW experiments within ±1.3 K, regardless of the heating rate used in the reactor. Each (cHAP, T) data point obtained in the MZW experiments corresponds to the mean of three simultaneous readings performed by the turbidity system at 550, 600, and 650 nm, respectively. As expected a tendency for a MZW widening (up to a maximum of 5 K), due to the lowering of the crystallization temperature as the cooling rate increases, was observed. Four

Figure 7. Phase diagram for the HAP + H2O system. Data obtained with the CB1 reactor at 12 K·h−1 (○), 18 K·h−1 (□), and 24 K·h−1 (△); green solid diamonds, equilibrium solubility by the gravimetric method; point a (cHAP = 56 g·kg−1; Ts = 333 K), solubility determined in a separate experiment were the HAP + H2O mixture was progressively heated in a jacketed Schlenk cell until complete dissolution was visually observed. The concentration of HAP, cHAP, refers to grams of anhydrous solute per 1 kg of water.

main domains are noted in Figure 7: (i) homogeneous solution, in red; (ii) metastable zone, in green; (iii) colloidal phase, in yellow; and (iv) solid+solution, in blue. This last domain can be further divided into three subzones (I, II, and III), where five solid forms with different lifetimes could be identified: zone I, approximately corresponding to the concentration range 16 g·kg−1 < cHAP < 30 g·kg−1; zone II, for 30 g·kg−1 < cHAP < 37 g·kg−1; and zone III, for 37 g·kg−1 < cHAP < 60 g·kg−1. The homogeneous solution (red) domain is delimited by the liquidus line (AA′), which represents the temperatures, Ts, at which the HAP + H2O system becomes completely liquid on heating. This curve was delineated by combining the Ts data obtained by the gravimetric method and by using the CB1 reactor. The upper extreme of the AA′ line was defined based on point a (cHAP = 56 g·kg−1; Ts = 333 K) which was determined in a separate experiment were the HAP + H2O mixture was progressively heated in a jacketed Schlenk glass cell until complete dissolution was visually observed. This experi2937

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Figure 8. Images of an HAP + H2O mixture with cHAP = 56 g·kg−1: (a) homogeneous solution at 333 K; (b) colloidal dispersion at 323 K; and (c) solid suspension at 300 K.

Figure 9. (a) Typical turbidity profile recorded at 600 nm for solutions of 18 g·kg−1 (zone I), 37 g·kg−1 (zone II), and 50 g·kg−1 (zone III). (b) Temperature−time profile illustrating the observed shift when crystallization occurs at points A, C, and F.

ment was necessary because for heating runs where cHAP > 39 g·kg−1 the turbidity probe was blind to the disappearance of the colloidal phase (see Cooling Crystallization Studies in the Materials and Methods section). The metastable zone width (in green) is delimited by lines AA′ and BCD. On cooling solutions with cHAP < 30 g·kg−1 (line BC) supersaturation directly leads to the crystallization of a solid phase (zone I). Above this concentration, however, the appearance of HAP crystalline phases (zones II and III) was mediated by the formation of a colloidal dispersion (line CD). An image of such dispersion, obtained on cooling a 56 g·kg−1 HAP+H2O mixture, is compared in Figure 8 with images of the precursor homogeneous solution and final solid suspension. The process is reversible and the dispersion can also be generated by heating the solid suspension. Preliminary dynamic light scattering (DLS) experiments indicated that the dispersion was formed by particles with diameters in the range 100−800 nm. A full DLS characterization of the colloidal phase as a function of temperature and HAP concentration will be described elsewhere. Turbidity−Time and Temperature−Time Profiles. Evidence that dissimilar crystallization processes might be occurring in the concentration ranges of zones I−III was first provided by the analysis of the turbidity− and temperature−time profiles obtained on cooling HAP + H2O solutions of different concentrations in the CB1 reactor (Figure 9). As shown in Figure 9a the turbidity profiles (normalized both in terms of

time and intensity so that the corresponding maxima equal unity) characteristic of zones I, II, or III are clearly different. On cooling (left to right direction in Figure 9a) solutions with HAP concentrations within zone I, only an abrupt increase of dispersed light intensity was noted (point A in Figure 9a) at the crystallization onset. This was accompanied by an inflection in the corresponding temperature−time profile (Figure 9b), due to the exothermicity of the crystallization process. When a similar run was carried out in zone II, a small step in the dispersed light intensity was first noted at the onset of the colloidal dispersion formation (point B). This event, which had no signature in the temperature−time profile, was followed by a progressive increase of the turbidity signal up to point C where crystallization occurred. The transition from colloidal dispersion to crystal formation was marked by a negative peak in the turbidity curve between points C and D. This occurrence was also signaled by an inflection in the temperature−time curve, analogous to that illustrated in Figure 9b. After the peak only a slight progressive increase of the dispersed light intensity was noted. Finally, in the case of solutions with concentrations corresponding to zone III, the turbidity−time profile first showed a small step at point E (similar to that noted for zone II mixtures) indicating the appearance of the colloidal dispersion. A gradual increase of the turbidity signal was then observed up to point F, where a sudden intensity drop marked the onset of crystallization. From then on the signal exhibited a constant 2938

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Figure 10. Images of the HAP crystals obtained on cooling (18 K·h−1) solutions with concentrations of (a) 17 g·kg−1 (zone I), (b and c) 32 g·kg−1 (zone II), and (d) 45 g·kg−1 (zone III). Panel (b) was captured immediately after the precipitation of the crystalline material, which rapidly (∼60 s) evolved into the product shown in panel (c).

(Figure 10c) after ∼60 s. These prisms were identified as the anhydrous form II HAP by XRPD (Figure 11). Indeed the corresponding powder pattern could be indexed as orthorrombic, space group P212121, a = 6.1406(26) Å, b = 9.5132(25) Å, c = 24.3443(49) Å (see Supporting Information for indexation details) in good agreement with analogous data from single crystal X-ray diffraction analysis (orthorhombic, space group P212121, a = 6.1097(11) Å, b = 9.5293(14) Å, c = 24.313(4)

positive deviation from the initial baseline (see inset at the bottom of Figure 9). Again, the temperature−time profile was sensitive to the crystallization event, but not to the formation of the colloidal dispersion. Characterization of Solid Phases. To elucidate the origin of the differences observed in the turbidity and temperature change patterns, the cooling crystallization of mixtures with cHAP = 17 g·kg−1 (zone I), cHAP = 32 g·kg−1 (zone II), and cHAP = 45 g·kg−1 (zone III) was visually monitored at a rate of 18 K·h−1, under magnetic stirring, in a 100 cm3 jacketed glass vessel (see Cooling Crystallization Studies on the Materials and Methods section). As mentioned above, this setup allowed sampling the solid phases for microscopy, XRPD, and TG analysis. Figure 10 shows optical microscopy images of crystals corresponding to zones I−III. The material obtained in zone I consisted of small needle-shaped crystals forming cotton-like aggregates (Figure 10a). These crystals were shown by XRPD to correspond to the above-mentioned H2 hydrate phase, whose powder pattern is illustrated in Figure 3. This was further supported by TG analysis which consistently gave the stoichiometry HAP·3H2O. The reproducibility of the TG results for products from different crystallization runs within zone I also suggested that the hydrate H2, and not a phase mixture, was formed. In zone II, crystallization initially led to very long and thin needles (Figure 10b) that were fully converted into prisms

Figure 11. Comparison of (a) the experimental X-ray powder diffractogram of the prismatic material obtained from cooling crystallization of a HAP solution with cHAP = 32 g·kg−1 (zone II) and (b) that of form II HAP simulated from published6 single crystal X-ray diffraction results. Both patterns refer to 296 ± 3 K. 2939

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Å).6 The lifetime of the needle-shaped form could be considerably extended (∼7 days) if the crystallization experiment was performed without stirring. It was also found that when collected from the mother liquor the crystals quickly (∼5 min) evolved into an opaque crystalline solid identified by XRPD as form I HAP (and not form II as when in contact with the solution). This transformation suggested that the original needle-shaped compound might be a hydrate form. Moreover, the fact that the transformation was too fast to allow a reliable XRPD characterization indicated that the material possibly corresponded to a third hydrate phase (H3) distinct from the hydrates H1 and H2. As shown in Figure 12, the recorded Figure 13. Comparison of (a) the experimental X-ray powder diffractogram of the prismatic crystals obtained from cooling crystallization of a HAP solution with cHAP = 45 g·kg−1 (zone III) and (b) that of form I HAP simulated from published6 single crystal Xray diffraction results. Both patterns refer to 296 ± 3 K.

HAP hydrates was never observed in zone III. Nevertheless, because in this zone the crystallization temperature is always close or below 302 K (the frontier of the H2/form I stability domains, see Figures 1 and 7) the form I → H2 transformation subsequently occurred within hours (cf. Scheme 1).

4. CONCLUSIONS The results here reported showed that five different HAP crystalline phases (the previously characterized anhydrous forms I and II6 and the hydrate H1,7 and two new hydrates H2 and H3) can be selectively and reproducibly obtained by cooling crystallization from water, provided that the experimental conditions are judiciously chosen based on the experimentally obtained T−cHAP phase diagram (Figure 7). Preliminary slurry tests carried out in ethanol confirmed the enantiotropic relationship between the anhydrous forms, but suggested that the equilibrium temperature of the form II → form I transition may be ∼50 K lower than judged from previous DSC results. The increase of the cooling rate from 12 K·h−1 to 24 K·h−1 led to an enlargement of the MZW in water up to a maximum of 5 K. This change did not, however, affect the nature of the crystallized materials. The crystallization precedence of the various phases, their approximate lifetimes, and transformation sequences while in contact with the solution are summarized in Scheme 2. On cooling solutions within the concentration range 16 g·kg−1 < cHAP < 30 g·kg−1 (zone I in Figure 7) the hydrate H2 of stoichiometry HAP·3H2O first crystallizes. For HAP concentrations in the range 30 g·kg−1 < cHAP < 37 g·kg−1 (zone II), crystallization initially leads to a very unstable hydrate phase (H3), whose stoichiometry could not be determined. If kept in contact with the solution under stirring, the H3 phase evolved into form II HAP after a few seconds. For HAP concentrations in the range 37 g·kg−1 < cHAP < 60 g·kg−1 (zone III) the crystallization of form I HAP is observed. Because the crystallization of all the above-mentioned phases occurs below ∼302 K (the upper temperature limit of the H2 domain in Figure 1) they all evolve into the H2 hydrate (HAP·3H2O) in a few hours. This phase is, however, found to always transform into the lower hydrate H1 (HAP·1.5H2O) over a period of months at least at temperatures close to 293 K. It is also interesting to point out that the cooling crystallization of HAP solid phases from solutions with cHAP

Figure 12. Overlay of the experimental X-ray powder diffractogram of the unstable needle-shaped crystals (H3) obtained from cooling crystallization of a HAP solution with cHAP = 32 g·kg−1 (zone II) and those of hydrates H1 and H2 and that of the anhydrous form I HAP simulated from published6 single crystal X-ray diffraction results. All patterns refer to 296 ± 3 K. The main peaks from H1 and H2 that are not present in the pattern of H3 are marked with an asterisk.

diffractogram coincides to a great extent with that of form I HAP, indicating that the initial sample had undergone significant transformation during the analysis. There is, however, a peak at 2θ = 8.75° (hence, acquired at the beginning of the experiment), that does not belong to the previously characterized anhydrous forms I and II,6 but which, as shown in Figure 12, has equivalents in the case of the hydrates H1 (2θ = 8.78°) and H2 (2θ = 8.87°). Nevertheless, the fact that (i) no decomposition H1 or H2 was observed during the recording of their XRPD patterns and (ii) several peaks typical of H1 and H2 (marked with an asterisk in Figure 12) are absent in the pattern of H3, suggests that the needlelike crystals obtained in zone II correspond, indeed, to a different hydrate and not to a mixture of form I with either H2 or H3. The H3 form was, however, too unstable to allow a stoichiometry determination by TG or a structure analysis by single crystal X-ray diffraction. Finally the crystals from zone III (Figure 10d) were found to correspond to the anhydrous HAP form I (in spite of their similarity with the form II crystals obtained in zone II, see Figure 10c). This was evidenced by the indexation of their XRPD pattern (Figure 13, see Supporting Information for indexation details), which led to monoclinic, space group P21/c, a = 7.7070(28) Å, b = 8.3336(31) Å, c = 11.2911(36) Å, β = 94.91(3)°. These results are, indeed, in good agreement with those obtained from a previous single crystal X-ray diffraction determination on form I HAP, namely monoclinic, space group P21/c, a = 7.7200(15) Å, b = 8.3600(17) Å, c = 11.280(2) Å, and β = 95.02(3)°.6 Note also that the direct crystallization of 2940

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Scheme 2

> 30 g·kg−1 was mediated by a colloidal dispersion, which was also observed when HAP slurries were heated above 323 K. Preliminary dynamic light scattering experiments indicated that this dispersion, whose stability domain could be defined in the T−cHAP phase diagram (Figure 7), consisted of particles with diameters in the range 100−800 nm. Finally, the present study highlights the unequivocal value of phase diagrams in the control of the selective and reproducible production of different solid forms by cooling crystallization. It also illustrates some of the complexities arising from the neverending wrestle of thermodynamics with kinetics in determining the phase that first crystallizes from solution and its stability over time, while in contact with or when removed from the mother liquor.

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(3) Bernstein, J. Polymorphism in Molecular Crystals; Oxford University Press: Oxford, 2002. (4) Hilfiker, R. Polymorphism in the Pharmaceutical Industry; WileyVCH Verlag GmbH & Co.: Weinheim, 2006. (5) Mullin, J. W. Crystallization, 4th ed.; Butterworth-Heinemann: Oxford, 2001. (6) Bernardes, C. E. S.; Piedade, M. F. M.; Minas da Piedade, M. E. Cryst. Growth Des. 2008, 8, 2419−2430. (7) Bernardes, C. E. S.; Piedade, M. F. M.; Minas da Piedade, M. E. Cryst. Growth Des. 2010, 10, 3070−3076. (8) Mangin, D.; Puel, F.; Veesler, S. Org. Process Res. Dev. 2009, 13, 1241−1253. (9) Oonk, H. A. J.; Calvet, M. T. Equilibrium Between Phases of Matter; Springer: Dordrecht, 2008. (10) Laugier, J.; Bochu, B. Checkcell; http://www.ccp14.ac.uk/ tutorial/Imgp. (11) Macrae, C. F.; Bruno, I. J.; Chisholm, J. A.; Edgington, P. R.; McCabe, P.; Pidcock, E.; Rodriguez-Monge, L.; Taylor, R.; van de Streek, J.; Wood, P. A. J. Appl. Crystallogr. 2008, 41, 466−470. (12) Bernardes, C. E. S.; Minas da Piedade, M. E. J. Therm. Anal. Calorim. 2010, 100, 493−500. (13) Gonçalves, E. M.; Minas da Piedade, M. E. J. Chem. Thermodyn. 2012, 47, 362−371. (14) Burger, A.; Ramberger, R. Mikrochim. Acta 1979, 2, 259−271. (15) Threlfall, T. L. Org. Process Res. Dev. 2009, 13, 1224−1230.

ASSOCIATED CONTENT

S Supporting Information *

Tables S1−S3 with detailed results of the solubility, MZW, and TG measurements. Tables S4−S7 with the indexation data of Figures 4, 5, 11a, and 13a. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by Fundaçaõ para a Ciência e a Tecnologia, Portugal through Projects PTDC/QUI-QUI/ 098216/2008, PEst-OE/QUI/UI0612/2011, and PEst-OE/ QUI/UI0100/2011. A postdoctoral grant (SFRH/BPD/ 43346/2008) awarded to C.E.S.B. is also gratefully acknowledged.

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REFERENCES

(1) Brittain, H. G. Polymorphism in Pharmaceutical Solids; Marcel Dekker: New York, 1999. (2) Brittain, H. G. Polymorphism in Pharmaceutical Solids, 2nd ed.; Informa Healthcare: New York, 2009. 2941

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