Tu Lee, PhD

)-(±)-2-(4-isobutylphenyl) propionic acid is a popular active pharmaceutical ingredient (API) in analgesic, anti-inflammatory, and antipyretic therapies (1) with abundant information about the compound provided in the literature (2–11). Ibuprofen is marketed as a racemate. The enantiomer (

)-(+)-ibuprofen is the active agent and the other enantiomer, (R)-(–)-ibuprofen, is partially converted into (

)-(±)-ibuprofen has several disadvantageous formulation properties such as poor water solubility of less than 1 mg/mL at 25 °C, a low melting point of 77 °C, and possible esterification with excipients containing a hydroxyl group (1). These problems can be easily overcome by using an (

)-(±)-ibuprofen salt such as the racemic compound (

)-(±)-sodium 2-(4-isobutyl–phenyl) propionate dihydrate or (

)-(±)-sodium ibuprofen dihydrate (see Figure 1), whose potential values over (

)-(±)-ibuprofen make the study of its solubility, polymorphism, crystallinity, crystal habit, and drying scheme meaningful (12).  

-(±)-sodium ibuprofen dihydrate solids at 25 °C was determined in 23 solvents using a robust and scalable initial solvent-screening method (13). Good solvents were identified, and the solubility curves of the solids in various good solvents were constructed. 

were defined as solvents in which the solubility of racemic  (

)-(±)-sodium ibuprofen dihydrate had a solubility of < 1 mg/mL at 25 °C. All solids were produced from the supersaturated solution of the good solvents by temperature cooling from 60–25 °C.  

Differential scanning calorimetry (DSC), thermogravimetric analysis (TGA), and powder X-ray diffraction (PXRD) determined the polymorphism, crystallinity, and drying scheme of the solid samples (14). Optical microscopy (OM) and scanning electron microscopy (SEM) were used for crystal-habit imaging. The solids' solubilities and crystal habits were related to the solvent microscopic properties of dispersion forces (δ

) by the Hansen model adopted from the paint industry at 25 °C (15).  

Solvent liquids generally are held together by van der Waals forces, which are electromagnetic interactions among molecules. The strength of van der Waals forces is directly reflected by the heat of vaporization (Δ

 can be translated into a correlation between vaporization and the solubility behavior because the types of intermolecular attractive forces needed to be overcome to vaporize a solvent liquid are the same as those that must be overcome to dissolve it. The Hildebrand numerical value (δ

 ]) of the solvency behavior of a specific solvent was invented as a function of Δ

in which R is the gas constant (8.314 J/mol • K), 

Hansen parameters fine-tune the total Hildebrand value (δ

) into three contributions: a dispersion-force component  (δ

) (15) as expressed in the following equation:  

are the induced attractions by the random polarities of two molecules in proximity. They are called

 London dispersion forces or induced dipole-induced dipole forces, 

and the number of temporary dipoles increases with the molecule's size. Polar forces are created when electrons are unequally shared between the individual atoms in a molecule. The charge distribution is related to the atomic composition, geometry, and size of a molecule. When polar molecules arrange themselves head-to-tail and positive-to-negative, the molecular network leads to a further increase in intermolecular interaction. These temperature-dependent dipole–dipole forces collectively create the Keesom orientation effect. A particularly strong type of polar interaction occurs when hydrogen's sole electron is drawn toward such electronegative atoms as oxygen, nitrogen, and fluorine, leaving the positively charged hydrogen nucleus exposed. The exposed positive nucleus exerts a considerable attraction on electrons in other molecules, forming a proton bridge, which is called 

Table I lists the 23 solvents from five different classes in the ascending order of Hildebrand parameters (13).  Two other solvents 

OH, ACS grade, 99.98%, boiling point of 97 °C, molecular weight of  60.1 g/mole, Lot 909065 Tedia Company, Fairfield, NJ) and cyclohexane (C

, ACS grade, 99.9%, boiling point of 80.7 °C,  molecular weight of 84.16 g/mol, Lot k24470166, Merck KGaA, Darmstadt, Germany) were used to verify the solubility sphere in the three-dimensional (3-D) Hansen plot.  

Table I: Total Hildebrand value and Hansen parameters for solvents at 25 8C.			

, molecular weight of 228.29 g/mol, batch number 015K0586) purchased from Sigma-Aldrich Corporation (St. Louis, MO) was a dihydrate of racemic (

)-(±)-sodium 2-(4-isobutylphenyl) propionate with a molecular formula of C

O and a molecular weight of 264.29 g/mol.  The amount of the stoichiometric water was determined by the thermogravimetric-use test showing about 13% water loss by weight near the boiling point of water (see Figure 2). 

)-(±)-sodium ibuprofen dihydrate, its solubility was determined by the gravimetric titration method that required a relatively small amount of material (13). Although the gravimetric titration method had an inherent error of about ± 20% because the solvent volume used was assumed to be the same as the solution volume (13) and was less accurate than the well-known UV-visible spectroscopic technique (16) and the evaporation-weighing technique (17), it was rapid, robust, and more suited for the initial screening. Unlike the other two techniques, the gravimetric titration method did not require establishing time-consuming calibration curves of (

)-(±)-sodium ibuprofen dihydrate concentration in every solvent and did not produce any unforeseen solvates that could interfere with the true mass of the dihydrate solids after removal of the  solvents.  

 The miscibility tests were performed by mixing two solvents from the 23 solvents in a 20-mL scintillation vial (13). The number of miscible and immiscible solvent pairs from the miscibility investigations and the API solubility in the 23 single solvents from the 

is defined as the total number of solid-generation experiments in pure-solvent, cosolvent and antisolvent systems (13). 

Polymorph, crystallinity, and crystal-habit studies.

 Solids were produced in the good solvents ferreted out by the 

 (13). The polymorphism, crystallinity, and crystal habit of the solids were determined by DSC, TGA,  PXRD, OM, and SEM, respectively. 

)-(±)-sodium ibuprofen dihydrate compound grown from water were filtered, rinsed with acetone, and vacuumed-dried at 40 °C or 4 h. The anhydrous (

)-(±)-sodium ibuprofen solids were mounted on an aluminum stub by a double-sided carbon conductive adhesive tape (Product number 16073,  Ted Pella Inc., Redding, CA).  The solid sample was sputter-coated with a 6-nm thick gold film in an Hitachi E-1010 Auto Sputter Coater (Hitachi Ltd., Tokyo, Japan). 

SEM was carried out using an Hitachi S-3500N (Hitachi Ltd., Tokyo, Japan) instrument equipped with a tungsten filament cathode source. Gold-coated samples were examined with beam energies of 15 kV with a chamber pressure of 10

  Pa (resolution ~3Å at these voltages). 

)-(±)-sodium ibuprofen dihydrate compound were dried by three various modes in a DOV-40 oven (Dengyng Instrument Co., Ltd., Taiepi, Taiwan) at 40 °C for 4 h in air, at 40 °C for 4 h under vacuum, and 90 °C  for 2 h under vacuum. A vacuum of 6.7 × 10

  Pa was generated by the GVD-050A vacuum pump (Ulvac Kiko Inc., Yokohama, Japan). 

)-(±)-sodium ibuprofen dihydrate dissolved well in nine good solvents in Classes 4 and 5: tetrahydrofuran (THF), 

-butyl alcohol, isopropyl alcohol (IPA), benzyl alcohol, 

-dimethylformamide (DMF), ethanol, dimethyl sulfoxide (DSMO), methanol, and water. The compound only slightly dissolved in 

-butyl ether, benzene, methyl ethyl ketone, chloroform, 

-dimethylaniline, acetone, 1,4-dioxane, nitrobenzene, and acetonitrile with a solubility of <1 mg/mL at 25 °C, giving 14 

solvents in Classes 1, 2, and 3. Only nine solubility curves of 

)-(±)-sodium ibuprofen dihydrate in nine good solvents at 15, 25, 40, and 60 °C were constructed and grouped by their similar solubility ranges for ease of comparison (see Figure 3). 

If the solute and the solvent formed an ideal solution, a straight line should result when 

 according to the van't Hoff equation (13): 

)-(±)-sodium ibuprofen dihydrate in the solution at a given temperature (

) measured in K, converted from the solubility curves in Figure 3, and 

 was the ideal gas enthalpy constant (8.314 J/mol • K). Any deviation from the ideal solution behavior would be reflected from the correlation coefficient (

) of the linear fit in Figure 4. The relatively high correlation coefficients with values > 0.97 of the solvent systems, except for the one of DMSO, indicated that the solubility measurements by gravimetric titration were quite accurate, and the deviation of the solution from ideality was negligible. The enthalpy of dissolution (Δ

-intercept of the straight lines in Figure 4, respectively. 

)-(±)-sodium ibuprofen dihydrate are summarized in Table II for qualitative comparisons. The positive values of Δ

 in Table II indicate that the energy of attraction of (

)-(±)-sodium ibuprofen dihydrate with each other, and the energy of attraction of the solvent molecules with each other were lower than the energy of attraction of (

)-(±)-sodium ibuprofen dihydrate and the solvent molecules in the solution. Heat was absorbed to make (

)-(±)-sodium ibuprofen dihydrate dissolve in the solvent. The solubility of (

)-(±)-sodium ibuprofen dihydrate increased with temperature.  The positive values of (Δ

)-(±)-sodium ibuprofen dihydrate-solvent systems became less ordered as (

)-(±)-sodium ibuprofen dihydrate dissolved into the solution. The entropy gain of the whole system was the main driving force for dissolution. 

Table II: Total Hildebrand values for solvents at 25 8C  versus enthalpies and entropies of solution of (R, S)-(6)-sodium ibuprofen dihydrate.			

Because of the symmetrical nature of the solvent miscibility table (see Table III), the number of boxes was divided by 2, except for the diagonal boxes (13). Based on the solvent miscibility studies of the solvent pairs of the 23 solvents, 36 gray boxes were divided by 2 to give 18 immiscible pairs in total (see Table III). The form space (

. a possible location for discovering a new polymorph) of the pure-solvent systems for the initial solvent-screening was limited to the number of good pure solvents (repesented by the yellow boxes). The form space for (

)-(±)-sodium ibuprofen dihydrate was 9. However, if the good cosolvent systems (

, binary miscible mixtures of good solvents) were taken into account, the form space would be extended to the total number of blue boxes in the solvent miscibility table divided by 2 (see Table III), which was equal to 36. In addition, if the antisolvent systems (

 binary miscible mixtures of a good and a bad solvent) also were considered, the form space of the antisolvent systems was calculated as the number of green boxes in the solvent miscibility table divided by 2 (see Table III); this value was 109. 

Table III. Solvent miscibility table, cosolvent, and antisolvent systems of (R,S)-(6)-sodium ibuprofen dihydrate. 			

Consequently, the total form space should then be at least equal to 9 + 36 + 109 =154. The total form space is expected to expand dramatically if various solvent compositions of binary mixtures, temperatures, and ternary solvent systems also are considered. Solid generation by temperature cooling only was applied in the yellow regions, yet the same crystallization mode could be applied to the blue regions in the solvent-miscibility table (see Table III). For the green regions (see Table III), solid generation is achieved isothermally by adding an antisolvent. Generally, no attempts are made for solid generation in the regions of immiscible solvent pairs (represented by the purple boxes), bad solvents (represented by the red boxes), and cosolvents of bad solvents (

., binary mixture of miscible bad solvents) (represented by the white boxes) in the solvent miscibility table (see Table III). References of other miscible solvents pairs from organic solvents other than the 23 solvents listed in Table III are available in the literature (19, 20).  

When the positions of the 23 solvents were located within  the 3-D space of a dispersion-force component (δ

) according to their corresponding coordinates of (δ

) in Table I and all the good solvents were represented by yellow and bad solvents by red in the form space (see Table III), a cluster of yellow domains were formed. The contour of these yellow domains outlined a 3-D volume of solubility in space. Good solvents were those solvents within the volume, bad solvents were solvents outside the volume (see Figure 5). This space could be represented by a solubility sphere with center coordinates (δ

)-(±)-sodium ibuprofen dihydrate in the 3-D Hansen plot. 

By trial and error, the center coordinates (δ

) were determined using the following (15): 

 was the distance between the solvent location and the center of the sphere. For the  23 solvents listed in Table I with their corresponding sets of Hansen parameters (δ

 would result for any chosen set of center coordinates (δ

). Only one particular set of center coordinates of (δ

 values for nine corresponding good solvents that were < the 14 D

 calculated from the 14 bad solvents. The largest value among the nine D

  values was the interaction radius of the solubility sphere.  For (

)-(±)-sodium ibuprofen dihydrate at 25 °C, the values of δ

with respect to the 23 solvents listed in Table I were determined to be 19.89, 16.19, 30.89 and 25.92 MPa

-propanol and cyclohexane were calculated using Equation 4. They were 18.18 and 35.24 MPa

-propanol was inside the solubility sphere (18.18 MPa

 ), but cyclohexane was outside (35.24 MPa

 ). These calculations agreed with the experimental observations that (

)-(±)-sodium ibuprofen dihydrate dissolved well in 

-propanol at 25 °C with a solubility of 53 mg/mL but dissolved poorly in cyclohexane at 25 °C with a solubility < 0.27 mg/mL. 

Table IV: Enthalpy of dehydration, enthalpy of melting, and crystallinity of  (R, S)-(6)-sodium ibuprofen dihydrate produced from seven of nine solvents arranged by the total ascending Hildebrand values.			

Solids generated from all nine good solvents were isolated and analyzed with DCS, TGA, and PXRD. Solids, however,  grown from benzyl alcohol and DMSO (two solvents with high-boiling points) were still wet after oven-drying in air at 40 °C for 4 h and formed solvates with ill-defined DSC and TGA results. Only seven of nine solids could be represented by a typical DSC response of solids generated from water (see Figure 6). The wide-base endotherm from 50–100 °C  as supported by the boiling point of water and an approximately 13% weight loss in Figure 2 was corresponded to the enthalpy of dehydration. The enthalpies of dehydration (Δ

) of  seven solvents are summarized in Table IV. The enthalpies of dehydration were 6.6–12.2 kcal/mol of water loss (23).  

Based on the similar DSC response (see Figure  6) and weight-loss pattern (see Figure 2) for all solids grown from THF, n-butyl alcohol, IPA, DMF, ethanol, methanol,  and water, only one  pseudopolymorph was identified for (

)-(±)-sodium ibuprofen dihydrate. The typical DSC response (see Figure 6) showed that the chemically bonded water molecules sandwiched between a layer of all (

)-(±)-sodium ibuprofen dihydrate compound were completely removed near 100 °C (see Figure 2) (12).  

An endothermic solid–solid transformation with structural editing from a racemic compound to an anhydrous racemic conglomerate started to take place as indicated by the many small peaks from 100–190 °C because the enthalpies of fusion of 

 interactions (10.49 kcal/mol) were higher than those of 

 interactions (8.71 kcal/mol) (24).  As the temperature increased, the anhydrous (

)-(±)-sodium ibuprofen eventually melted into liquid around 200 °C . The sandwiched layer of water in the racemic (

)-(±)-sodium ibuprofen dihydrate also was completely removed by vacuum-drying either at 40 °C for 4 h or at  90 °C for 2 h (see Figure 7) (24). The DSC scan of an anhydrous sodium ibuprofen obtained by vacuum-drying either at 40 °C for 4 h or at  90 °C for 2 h exhibited a hump from 50–90 °C (see Figure 7), indicating an endothermic, ongoing diffusive restructuring of the 

 packing to form a higher energy state of the 

 packing that had not been completely finished during the annealing period at 40 °C for 4 h or 90 °C for 2 h under vacuum. 

)-(±)-ibuprofen dihydrate solids grown from water and oven-dried at 40 °C for 4 h in air is shown in Figure 8. A SEM micrograph (see Figure 9) showed when the water molecules of racemic (

)-(±)-sodium ibuprofen dihydrate partially were removed under the high vacuum inside the SEM chamber, the 

-layer stacked in the [001] direction were unzipped like a zipper in the [100] direction as seen in Figure 10, which was redrawn (12) by computer software (Diamond 3.1, Crystal Impact GbR, Brandenberg Germany). The authors believe the tiny elongated pores of the anhydrous sodium ibuprofen gave rise to its hygroscopic nature. 

Crystallinity of all crystals harvested from solvent screening was approximated using the following relationship (13, 25):  

All of the DSC endotherms (melting peaks) for racemic (

)-(±)-sodium ibuprofen dihydrate crystals produced from seven of nine good solvents and oven-dried at 40 °C for 4 h in air were compared with the DSC endotherm for racemic (

)-(±)-sodium ibuprofen dihydrate crystals grown from methanol with the highest heat of melting of 54.95 J/g near 200 °C as a standard. These results are summarized in Table IV. 

Articles

The racemic compound (R, S)-(±)-ibuprofen is a popular and well understood active pharmaceutical ingredient, but it has several disadvantageous formulation properties such as poor solubility, low melting point, and potential esterification with excipients containing an hydroxyl group. The authors investigate the use of an (R, S)-(±)-ibuprofen salt to evaluate these problems using various analytical methods to determine the polymorphism, crystallinity, and drying scheme.

Solubility, Polymorphism, Crystallinity, Crystal Habit, and Drying Scheme of (R, S)-(±)-Sodium Ibuprofen Dihydrate

Solubility, polymorphism, crystallinity, and crystal habit of an active pharmaceutical ingredient (API) play critical roles in the value chain of pharmaceutical development, manufacturing, and formulation (1–3). The solubility of an API in solvents and solvent mixtures has a considerable influence on the choice of solvents and the course of operation in solvent-based processes such as chemical reaction, extraction, crystallization, filter cake washing, and wet granulation (4–9). The polymorphism of an API determines its packing, thermodynamic, spectroscopic, kinetic, surface, and mechanical properties in the solid state (10). The crystallinity of an API contributes to the mechanical properties of a compact, an API's stability, and the dissolution rate (11–15). The crystal habit of an API also has profound effects on the rate at which the API can be processed in filtering, washing, and drying, and the success of the API in powder flow, blending, direct compaction, roller compaction, wet granulation, and dissolution rate (16–19). 

Because solubility, polymorphism, crystallinity, and crystal habit are all solvent dependent, solvent screening is of fundamental and foremost importance to many chemical process industries, especially the pharmaceutical industry (4, 20–22). Recently, there has been increased interest in performing high-throughput polymorph screening in miniaturized scales using solvent evaporation (23–25). Because of the course of solid generation from a supersaturated solution by evaporation and the relatively small amount of the API used, however, these methods may not always correlate directly with the scale-up conditions in crystallization (

 a relatively large volume of solvent, temperature cooling, and stirring) and do not provide direct information about other simultaneous effects brought about by the solvent such as an API's solubility, crystallinity, and crystal habit. 

This article promotes an alternative solvent-screening strategy that is tailor-made for the drug development and design of API solids processes. Under this initial screening strategy of pure-solvent systems, 23 kinds of solvent mostly useful for scale-up were chosen (26). The solubility of the API solute in each solvent at 15, 25, 40, and 60 °C was measured by gravimetric titration. The enthalpy and the entropy of solution for each solvent system were calculated. The solubility of the API in each solvent at 25 °C was plotted against the dielectric constant of various solvents, resulting in a characteristic solubility pattern that might serve as the fingerprint to identify a particular API. Although only pure-solvent systems were being considered in this study, the total "form space" for each API—that is, the total number of solid generation experiments in pure-solvent, cosolvent, and antisolvent systems— was also calculated on the basis of the number of good solvents for the API from the solubility studies and the number of miscible and immiscible solvent pairs from the miscibility investigations.  

Solid generation of the API solute in each pure solvent was achieved by gently shaking a 20-mL scintillation vial and by temperature cooling from 60 to 25 °C under an ambient condition (27). The cooling rate of a solution with a volume <20 mL was almost independent from the volume and the nature of solvent. The cooling profile could be approximated by a quadratic equation determined experimentally as 

 is the time (min). The relatively rapid decrease in temperature serves as an ideal way to induce a polymorph that normally does not occur thermodynamically. In addition, temperature cooling is a common method in crystallization scale-up. 

Differential scanning calorimetry (DSC) was used to determine the polymorphism and the crystallinity of the API solid samples (28). Optical microscopy (OM) was used for crystal habit imaging. Two well-known APIs were selected for our solvent screening strategies: acetaminophen (4-acetamidophenol) and racemic (±)-ibuprofen (α-methyl-4- [isobutyl] phenylacetic acid) (see Figure 1), because of their worldwide commercial values in analgesic and antipyretic therapy, their abundant information in the literature (4,11,17, 29–44), and the lack of extensive solvent studies on their solubility, polymorphism, crystallinity, and crystal habit (29–31). 

 Figure 1: Active pharmaceutical ingredient solute molecules.			

 The following solvents were used in this study: cetone (CH

, HPLC–Spectro grade, 99.5%, boiling point [bp] = 56 °C, molecular weight [MW] = 58.08, lot 411050), acetonitrile (CH

CN, ACS grade, 99.96%, bp = 81.6 °C, MW = 41.05, lot 812045), benzene (C

, ACS grade, 99%, bp = 80 °C, MW = 78.11, lot 310008), benzyl alcohol (C

OH, certified grade, 99.9%, bp = 205 °C, MW = 108.14, lot 809010), 

OH, ACS grade, 99.4%, bp = 117.7 °C, MW = 74.12, lot 205027), chloroform (CHCl

, HPLC–Spectro grade, 99.9%, bp = 60.5–61.5 °C, MW = 119.38, lot 401059), dimethyl sulfoxide (DMSO) ((CH

SO, HPLC–Spectro grade, 99.8%, bp = 189 °C, MW = 78.13, lot 202060), N,N-dimethylformamide (DMF) (HCON(CH

, ACS grade, 99.8%, bp = 153 °C, MW = 73.10, lot 020505), 

, HPLC– Spectro grade, 99.4%, bp = 98 °C, MW = 100.21, lot 111110), isopropyl alcohol (IPA) ((CH

CHOH, HPLC–Spectro grade, 99.8%, bp = 82.4 °C, MW = 60.1, lot. 310067), methanol (CH

OH, HPLC–Spectro grade, 99.9%, bp = 64.7 °C, MW = 32.04, lot. 411070), methyl tert-butyl ether (MTBE) ((CH

, certified grade, 99.9%, bp = 55.2 °C, MW = 88.15, lot 712032), methyl ethyl ketone (MEK) (C

, ACS grade, 99.6%, bp = 81.6 °C, MW = 72.11, lot 201021), tetrahydrofuran (THF) (C

O, HPLC–Spectro grade, 99%, bp = 65–67 °C , MW = 72.11, lot 411013), toluene (C

, HPLC–Spectro grade, 99.8%, bp = 110.6 °C, MW = 92.14, lot 406050), and xylene (C

, ACS grade, 98.5%, bp = 137–144 °C, MW =106.17, lot 305065). These solvents were obtained from Tedia Company (Fairfield, OH).  

The following were obtained from Acros Organics (Morris Plains, NJ): N,N-dimethylaniline (DMA) (C

, ACS grade, 99%, bp = 193 °C, MW = 121.18, lot  A0213203001), nitrobenzene (C

, ACS grade, 99%, bp = 210–211 °C, MW = 123.11, lot A019568301) and p-xylene (C

, ACS grade, 99%, bp = 138 °C , MW = 106.17 °C, lot 48754/2).  

The following were obtained from Showa Chemical Co., Ltd. (Tokyo, Japan): 1,4-dioxane (C

, ACS grade, 98%, bp = 100–102 °C, MW = 88.11, lot sp-3432R). Ethanol (CH

OH, HPLC–Spectro grade, 99.5%, bp = 78 °C, MW = 46.7 ) was obtained from Echo Chemical Co. Ltd. (Taipei, Taiwan). Ethyl acetate (CH

, ACS grade, 99.5%, bp = 76.5–77.5 °C, MW = 88.11, lot G43342) was obtained from Grand Chemical Co. Ltd. (Daejeon, South Korea). Reversible osmosis water was clarified with a water-purification system (Milli-RO Plus, Millipore, Billerica, MA). 

OH, MW = 151.17, lot S23272-444) and racemic (±)-ibuprofen (α-methyl-4-[isobutyl] phenylacetic acid, C

, MW = 206.3, lot 26H1368) were obtained from Sigma-Aldrich (Steinheim, Germany). 

 Approximately 10 mg of the API were weighted in a 10-mL scintillation vial. Drops of solvent were titrated carefully with a micropipette into the vial with intermittent shaking until all API solids were just dissolved. The solubility of the API at a given temperature was calculated as the weight of the API in a vial divided by the total volume of the solvent added to a vial. Solubility of the API in the same solvent at 15, 25, 40, and 60 °C was determined. All temperatures were maintained and controlled by a water bath. Despite the inherent inaccuracy of this method in volume measurements by eye, it provided a rapid and robust way for process scale-up. 

 Out of the 23 solvents, approximately 1-mL portions of each solvent comprising a pair were shaken together for ~1 min in a 20-mL scintillation vial. If no interfacial meniscus was observed after the contents of the vial were allowed to settle, the solvent pair was considered to be miscible. If a meniscus was observed without apparent change in the volume of either solvent, the pair was regarded as immiscible. 

Polymorph, crystallinity and crystal habit studies.

 Saturated API solutions of each solvent were prepared in a 20-mL scintillation vial in accordance with the solubility values at 60 °C that were determined from the experiments described in the previous section. Solids were generated by cooling from 60 to 25 °C in an ambient condition with intermittent shaking. Solids were filtered with a strong vacuum and vacuum-oven dried at 55 °C overnight. Polymorphism, crystallinity, and crystal habits were determined by DSC and optical microscopy, respectively. 

 A differential scanning calorimeter (DSC 7, Perkin Elmer, Norwalk, CT) was used to monitor the thermal events during heating. The instrument was calibrated with indium (8–10 mg, 99.999% pure, extrapolated melting onset at 156.6 °C). All samples were run in sealed aluminum pans under a constant nitrogen purge. Each sample was heated at 10 °C/min. 

 An optical microscope (SZII, Olympus, Tokyo, Japan) equipped with a CCD camera (SSC-DC50A, SONY, Tokyo, Japan) was used to take images of crystal habit. 

The solvents were categorized into five classes according to their functional group (26):								

Class 1: "protic" or hydrogen bond donating solvents (Lewis acids), including benzyl alcohol, 

-butyl alcohol, ethanol, IPA, methanol, and water;

Class 2: hydrogen bonding acceptor solvents (Lewis bases), including acetone, 1,4-dioxane, DMF, ethyl acetate, MEK, MTBE, nitrobenzene, THF, and water;

Class 3: polar aprotic solvents, better termed "nonhydroxylic solvents," including acetonitrile, DMA, DMF, and DMSO;

Class 4: chlorocarbon solvents, which include chloroform; 

C : hydrocarbon solvents, both saturated and unsaturated, including benzene, 

Acetaminophen dissolved well in 15 of the solvents (labeled as "good solvents") and only slightly dissolved in 8 of the solvents (labeled as "bad solvents") with a solubility <1 mg/mL at 25 °C. The bad solvents for acetaminophen were mainly solvents of Classes 2, 4, and 5 and included MTBE, DMA, chloroform, benzene, 

-heptane, toluene, p-xylene, and xylene. Therefore, only 15 solubility curves of acetaminophen in 15 good solvents of Classes 1 and 3 at 15, 25, 40, and 60 °C (50 °C was used in cases in which a solvent's boiling point is close to 60 °C) were constructed and grouped by their similar solubility ranges for comparison (see Figure 2). As for racemic (±)-ibuprofen, all classes of solvents were good solvents except for water at 25 °C. Therefore, solubility curves of racemic (±)-ibuprofen in 22 good solvents at 15, 25, 40, and 60 °C (either 50 or 55 °C was used in cases in which a solvent's boiling point is close to 60 °C) were constructed and grouped by their similar solubility ranges as well (see Figure 3). Using scattered published results (29, 45) to check  the accuracy of our data in Figures 2 and 3, the error percentage was within ±20%. 

Figure 2: Solubility curves of acetaminophen in 15 kinds of solvent.			

If the solute and the solvent were assumed to form an ideal solution, the solubility could be predicted from the van't Hoff equation: 

 is the mole fraction of the API solute in the solution, 

 is the gas enthalpy constant (8.314 J/mol K), Δ

 is the entropy of dissolution (46). Extracting 

 from each solubility curve and performing linear-curve fitting to ln 

 of the API in each solvent could be approximated from the slope and intercept of a straight line. The values of Δ

 for acetaminophen and racemic (±)-ibuprofen are summarized in Tables I and II.  

Table I: Dielectric constants of solvents versus enthalpies and entropies of acetaminophen and racemic (Â±)-ibuprofen solutions.			

 in Table I indicate that the energy of attraction of the API solute molecules with each other and the energy of attraction of the solvent molecules with each other are lower than the energy of attraction of the API solute and the solvent molecules in the solution. Therefore, heat was absorbed to make the API solute dissolve in the solvent. In this case, the solubility of the API solute increased with temperature. Furthermore, the positive values of Δ

 in Table I showed that the API solute–solvent systems became less ordered as the API solute dissolved into the solution. 

Table II: Solvent miscibility table, cosolvent and antisolvent systems of acetaminophen.			

When the solubility values of acetaminophen and racemic (±)-ibuprofen at 25 °C were plotted against the dielectric constant of various solvents (see Table I) in which the APIs were soluble, the APIs' characteristic "solubility pattern" was produced (see Figures 4 and 5). Nonetheless, there were no direct correlations between the dielectric constants of the solvents and the APIs' solubility values (see Figures 4 and 5) or between the dielectric constants of the solvents and the APIs' Δ

 values, nor between the dielectric constants of the solvents and the Δ

Figure 3: Solubility curves of racemic (Â±)-ibuprofen in 22 kinds of solvent.			

The implication is that the solubility was profoundly influenced by various intermolecular interactions. Besides the Coulombic (ion–ion or ion–dipole) interactions, other effects such as the Keesom force (rotating dipole–dipole force), the Debye force (dipole–induced-dipole force), the London force (induced-dipole–induced-dipole force), and the hydrogen bonding should all be considered. The high solvation power of the two "polar aprotic solvents"—DMF and DMSO—in the two solubility patterns (see Figures 3 and 4) suggested that solid generation might be difficult to carry out in extremely good solvents. 

Figure 4: Solubility spectrum of acetaminophen.			

Based on the solvent miscibility studies of the solvent pairs of the 23 kinds of solvent, there were 18 immiscible pairs (36 gray boxes divided by 2) (see Tables II and III). It was necessary to divide the number of boxes by 2 when a particular solvent 

. Again, there were no direct correlations between solvent dielectric constants (see Table I) and solvent miscibility values (see Tables II and III). For example, 1,4-dioxane with a dielectric constant of 2.2 was miscible with all other solvents in Tables II and III, but xylene (dielectric constant 2.0) was immiscible with DMF, DMSO, and water. Similar to solute–solvent solubility, the solvent–solvent miscibility also was profoundly influenced by the many aforementioned intermolecular interactions. 

Table III: Solvent miscibility table, cosolvent and antisolvent systems of racemic (Â±)-ibuprofen.			

The pure-solvent systems were represented as the 23 diagonal boxes of the solvent pair of 

 in the solvent miscibility tables (see Tables II and III). The "form space" of the pure-solvent systems for our initial solvent screening was limited to the number of good pure solvents (

 indicated by the yellow boxes in Tables II and III) for each API. Therefore, the form space for acetaminophen and racemic (±)-ibuprofen was 15 and 22 respectively (see Tables II and III). If the good co-solvent systems (

 the binary mixtures of good solvents) were taken into account, however, the form space of the good cosolvent systems was extended to the number of blue boxes in the solvent miscibility table divided by 2 (see Tables II and III). Therefore, the form space under these particular conditions for acetaminophen and racemic (±)-ibuprofen was 102 and 224 respectively (see Tables II and III).  

Table IV: Form space of acetaminophen and  racemic (Â±)-ibuprofen.			

In addition, if the antisolvent systems (

 binary mixtures of a good and a bad solvent) also were considered, the form space of the antisolvent systems was calculated as the number of green boxes in the solvent miscibility table divided by 2 (see Tables II and III). The form space under these circumstances for acetaminophen and ibuprofen was 105 and 11, respectively (see Tables II and III) and the total form space for acetaminophen and racemic (±)-ibuprofen should then be at least equal to 222 and 257, respectively (see Table IV). 

Figure 5: Solubility spectrum of racemic (Â±)-ibuprofen.			

The total form space should expand dramatically if various solvent compositions of binary mixtures, temperatures, and ternary solvent systems are taken into account (23). Solid generation by temperature cooling can be applied in the yellow and blue regions in the solvent miscibility table (see Tables II and III). Nonetheless, a constant temperature must be maintained if solid generation was achieved by the addition of an antisolvent in the green domain (see Tables II and III). In general, no attempts are made for solid generation in the regions of immiscible solvent pairs (gray boxes), bad solvents (orange boxes), and cosolvents of bad solvents (

 binary mixtures of any two bad solvents) in the solvent miscibility table (see Tables II and III). In addition, references of other miscible solvent pairs from organic solvents other than the 23 kinds of solvent listed in Tables II and III are available in the literature (48, 49). 

Figure 6: Differential scanning calorimetry thermogram of a typical sample of acetaminophen.			

Out of the 15 good solvents, acetaminophen solids were generated in only 14 of them. N,N-dimethylformamide was so good a solvent that acetaminophen was unable to crystallize out by temperature cooling. All 14 isolated white solids were analyzed by DSC, resulting in a sharp endotherm at ~171 °C, which is the melting point of Form I crystals (see Figure 6) (37). This result clearly verified that acetaminophen powders solidified from all 14 solvents were the thermodynamically stable Form I crystals at 25 °C. As for the racemic (±)-ibuprofen, the white powders could only be generated from 18 out of the 22 good solvents. Racemic (±)-ibuprofen supersaturated solutions of acetone, benzyl alcohol, dimethyl sulfoxide, and xylene were unable to produce solids by temperature cooling. Nonetheless, the 18 isolated solid samples showed a typical endotherm around the solid–liquid melting point of 77 °C (see Figure 7) (30, 32). Apparently, there was no new polymorph and the fact that the racemic (±)-ibuprofen is isomorphic once again was verified (18). 

Solubility, polymorphism, crystallinity, and crystal habit of acetaminophen and racemic (+/-)-ibuprofen were determined by initial screening of 23 solvents for scale-up. Solubility curves were constructed, and solubility at 25 degrees Celsius was plotted against the dielectric constants of various solvent as a fingerprint for solute identification. The total "form space" for acetaminophen and racemic (+/-)-ibuprofen were calculated to be 222 and 257, respectively. Various crystal habits and sizes for ibuprofen and acetaminophen were observed.