Structure, Solubility, Screening, and Synthesis of Molecular Salts

Structure, Solubility, Screening, and Synthesis of Molecular Salts SIMON N. BLACK,1 EDWIN A. COLLIER,2 ROGER J. DAVEY,2 RON J. ROBERTS3 1

Process Engineering Group, PR&D, AstraZeneca, Macclesfield, Cheshire, SK10 2NA, United Kingdom

2 Molecular Materials Centre, School of Chemical Engineering and Analytical Science, The University of Manchester, Manchester, M60 1QD, United Kingdom 3

Preformulation & Biopharmaceutics, PAR&D, AstraZeneca, Macclesfield, Cheshire, SK10 2NA, United Kingdom

Received 28 July 2006; revised 25 October 2006; accepted 29 October 2006 Published online in Wiley InterScience (www.interscience.wiley.com). DOI 10.1002/jps.20927

ABSTRACT: The preparation of molecular salts as potential delivery vehicles for pharmaceutically active compounds is more common than current appreciation of the phenomena governing the solubility and isolation of salts suggests. In addition, it would appear that there are no reported measurements on a large enough data set for a serious structure–property relationship analysis to have been performed for this class of material. This means that at present, the ability to predict which salt forms will have desirable physical properties is essentially nonexistent. The work reported here sets out to explore these issues using new data on 17 salts obtained from a screen performed on the basic pharmaceutical ephedrine. The importance of solvent choice in salt formation, of salt selection in the control of bioavailability and of ternary phase equilibria in salt isolation and the relationship between a number of measured and calculated crystal properties are illustrated and discussed. The consequences of these relations for the general design, implementation, interpretation, and scale-up of salts screens are also explored. ß 2007 Wiley-Liss, Inc. and the American Pharmacists Association J Pharm Sci 96:1053–1068, 2007

Keywords: lization

crystal engineering; solubility; materials science; formulation; crystal-

INTRODUCTION The search for useful solid forms of pharmaceutically active molecules routinely includes a search for molecular salts.1 Such salts may offer advantages over the corresponding free acid or base in terms of physical properties such as melting point (thermal stability), crystallinity, hygroscopicity, dissolution rate, or solubility (bioavailability). Recently, there have been conEdwin A. Collier’s present address is Transform Pharmaceuticals, Hartwell Ave., Lexington, MA, USA. Correspondence to: Roger J. Davey (Telephone: 4401613064409; Fax: 4401613063909; E-mail: [email protected]) Journal of Pharmaceutical Sciences, Vol. 96, 1053–1068 (2007) ß 2007 Wiley-Liss, Inc. and the American Pharmacists Association

siderable technological advances that make it possible to carry out ‘‘salt screens’’ using very small amounts of active material together with a high degree of automation.2,3 Although these salt screens are common in the pharmaceutical industry, it is rare for the full results to be reported in the scientific literature. In particular, there is a lack of published data that include examples of both success and failure to give salts. The first aim of this study was therefore to use the total dataset from a salt screen to provide a fuller appreciation of the factors that determine success or failure in salt formation and hence, to provide a more scientific basis for salt screening protocols. The second objective was to provide a dataset of physical properties for a series of salts of a common

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base with a range of counter-ions. Such data could then be utilized to seek correlations between calculated crystal properties such as lattice energy and crystal density, and measured crystal properties such as temperature and enthalpy of melting and solubility. From a pharmaceutical viewpoint, the melting enthalpy, melting temperature, and solubility are of particular importance both because of their routine measurement and due to their potential influence on processing and bioavailability. There are surprisingly few previous studies of this type on molecular salts. In one early study,4 two correlations were proposed. The first was between the melting temperatures of salts of a basic molecule and the melting temperatures of the associated acids, while the second was between the melting temperatures of a range of basic drugs and the aqueous solubility of their salts. The new data set reported here was used to test these and other correlations and in the course of this study, a third aim emerged, namely to explore the phase diagrams of salt systems and to demonstrate their use in improving the scientific understanding of the—often problematic—scale-up procedures for salt syntheses.

Selection of a Model System For the work described in this study, the model pharmaceutical base (1R, 2S)-2-methylamino-1phenylpropan-1-ol, more commonly known as (1R, 2S)-(
)-ephedrine (hereafter ‘‘ephedrine’’) was selected. This choice was made on the basis that ephedrine, has a single pKa value, previously reported single crystal structure and analytical data, molecular weight <300, for ease of molecular modeling and is available at reasonable cost as a low melting point (37.58C), white crystalline solid.

Figure 1. (1R, 2S)-(
)-ephedrine salt structure.

The chemical structure of an ephedrine salt is shown in Figure 1, where A
represents the anion and the ephedrine cation is protonated at the secondary amine function NHþ 2 . Molecular modeling studies6,13 suggest that there are two stable conformations of the ephedrine molecule, only one of which contains an intramolecular hydrogen bond. Both conformations are observed in the available crystal structures of ephedrine and its salts.6–10 In previous studies, the lattice energies and physical properties of 11 diastereoisomeric salt pairs of ephedrine were compared in an attempt to rationalize the observed degree of enantiomeric separation by selective crystallization.6–8 The heats of fusions, melting points, and calculated lattice energies were reported, but not the aqueous solubilities. Additionally, Li et al.,14 within a study of 13 chloride and salicylate salts, included the hydrochloride salt of the single enantiomer of ephedrine selected here. None of these studies, however, provides a suitable dataset of the type required for our present purposes. For the 17 salts isolated and characterized in this current study, details of their crystal structures are reported separately15,16 and are available through the CSD.11

MOLECULAR CONSIDERATIONS

Ephedrine has a relatively simple molecular structure with a single pKa of 9.745 and a molecular weight of 165.24 g/mol. Crystal structures of the free base hemihydrate, hydrochloride salt, and selected phosphate and mandelate salts are available6–10 in the Cambridge Structural Database,11 but little has been documented concerning their crystallization conditions. Preparation of the 2-naphthalenesulfonate salt has been reported.12

The 25 anions used in this study (see Tab. 1) were chosen to give a distribution of charge and stereochemistry across the range of the five different classes of acidic counter-ions used regularly in the pharmaceutical industry. It is generally accepted that the most important molecular properties for the design of a salt screen are the acid and base dissociation constants (pKa).1 Using ephedrine acetate as an example, it is not difficult to see why. The speciation diagrams for ephedrine and acetic acid in water are shown in Figure 2a. In the pH range 6 to 8 both species exist predominantly as ions, favoring salt formation. This illustrates the basis for the oft-quoted ‘‘rule’’

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Previous Data on Ephedrine

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Table 1. The 25 Acids Used in This Study Carboxylic Acids Acetic Benzoic Formic Salicylic Trichloroacetic —

Dicarboxylic Acids

Hydroxy Acids

Inorganic Acids

Adipic Fumaric Maleic Malonic Succinic —

Citric D-(
)-gluconic Glycolic L-(
)-malic L-(þ)-tartaric —

Hydrochloric Nitric Phosphoric Sulfuric — —

Sulfonic Acids Benzene sulfonic 1,2-Ethane disulfonic Ethane sulfonic Methane sulfonic 2-Naphthalene sulfonic p-Toluene sulfonic

Figure 2. Speciation diagrams for ephedrine and acetic acid in (a) water–ephedrine pKa ¼ 9.74, acetic acid pKa ¼ 4.76 in; (b) methanol–ephedrine pKa ¼ 8.74, acetic acid pKa ¼ 9.71. DOI 10.1002/jps

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that for salt formation, the difference, DpKa, between the acid and base dissociation constants must be >2. In its application however, this rule is often used without appreciation of the potential impact of changing solvent on the value of DpKa. Figure 2b, for example, shows the impact on the speciation of changing the solvent from water to methanol. The pKa of ephedrine is not expected to change significantly, but the pKa of acetic acid shifts dramatically by 5 pH units.17 There is now no pH range for which both cation and anion co-exist, suggesting that it will be difficult to crystallize an acetate salt from methanol. All carboxylic acids used in this study have aqueous pKas in the range 2.8–5.7, and all will be shifted by 4–7 pH units in methanol. These values are shown in detail in Table 2. Overall this allows some predictions to be made about the outcome of salt screening experiments. In water, DpKa is large, so salt formation is expected for all counter-ions. For stronger acids, DpKa does not change significantly in methanol and hence, again salt formation is possible but for the weaker, carboxylic acids, DpKa is small or negative, so carboxylate salt formation will not be expected.

EXPERIMENTAL Crystallization Procedures Ephedrine base and the inorganic and organic acids were purchased from Sigma-Aldrich and used as supplied. The objective of crystallization experiments was to isolate sufficient crystalline material for characterization, as described below but including single crystal XRD, thermal and solubility measurements. In all cases, equimolar amounts of base and acid were required in a known volume of solvent, under conditions from which supersaturation could be generated either by cooling or evaporation. Thus crystallization experiments started either from liquors prepared

by mixing at room temperature, 0.1 M solutions of ephedrine and the various acids (evaporative crystallization) or from more concentrated equimolar solutions prepared at 508C by dissolving acid and base sequentially (cooling crystallization). In all cases, the pH of the starting solution was measured. Some evaporative crystallizations were performed in a highly solvent robust and optically inactive quartz, 96-well plate. Solutions were pipetted into each well using an autopipette, the plate was then covered with sealing tape, heated to 508C to ensure dissolution of all the solids and then uncovered to allow solvent evaporation. Analytical Methods Crystalline samples recovered from these experiments were characterized by a number of standard techniques: optical microscopy (Zeiss Axioplan 2 polarizing microscope linked to a color video camera), FTIR—Fourier Transformation infrared spectroscopy (Avatar 360 ESP integrated with Nicolet’s OMNIC software v.5.1b, fitted with an Attenuated Total Reflection (ATR) accessory), PXRD—powder X-ray diffraction (Bruker D8 Advance powder diffractometer with a sealed ˚ ) and positioncopper X-ray tube (l ¼ 1.5418 A sensitive detector), DSC—differential scanning calorimetry (Mettler Toledo DSC 820 with auto sampler and liquid nitrogen cooling controlled by a computer running Mettler Toledo’s STARe software v.6.10) and TGA—Thermogravimetric Analysis (Mettler Toledo TGA/SDTA 851e with auto sampler, controlled by a computer running Mettler Toledo’s STARe software v.6.10). Full details of these measurements and of some additional analyses are reported elsewhere.15 Solubility Determinations Once bulk samples of the ephedrine salt forms had been produced, the solubility of each was

Table 2. The pKas of Selected Acids in Water and Methanol (from Reference [11]) Compound

pKa Water

pKa Methanol

DpKa Water

DpKa Methanol

Ephedrine Acetic acid Benzoic acid Malonic acid Fumaric acid Succinic acid L-tartaric acid

9.74 4.76 4.19 2.83, 5.70 4.38 4.19, 5.61 3.02

8.69 9.63 9.30 7.66, 10.64 9.78 9.14, 11.30 8.12

— 4.98 5.55 6.91, 4.04 5.36 5.55, 4.13 6.72

—
0.94
0.61 1.03,
1.95
1.09
0.45,
2.61 0.57

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MOLECULAR SALTS

determined in water at 258C using a classic gravimetric method. Briefly, 5 mL of solvent was stirred overnight with excess solid in a jacketed vessel held at 258C. The pH of this suspension was measured and, on allowing the solids to settle, 2 mL of saturated solution was extracted with a syringe and dispensed through a 0.2 mm filter tip into a pre-weighed sample bottle. The sample was evaporated to dryness and solubility evaluated in the normal way. Solubility products were calculated from these data using a literature method;18 details of the calculations are given in reference 15. In the course of these calculations, a number of points became apparent which are worth reiterating here. First, the units of Ksp depend on the stoichiometry of the salt. For 1:1 salts of monoacids, (to within  3%) Ksp ¼ solubility2. For 1:1 salts of diacids, (except for adipate, malate, and tartrate monohydrate, where the measured pH is <1 unit from one of the pKa values) Ksp ¼ solubility2. For 2:1 salts, (except for edisylate, where the measured pH is <1 unit from the higher pKa), Ksp/4 ¼ solubility3. In these calculations, solubility is expressed as mols ephedrine/L. Lattice Energy Calculations CERIUS2 was used to calculate lattice energies for (1R, 2S)-(
)-ephedrine hemihydrate, (1R, 2S)(
)-ephedrine and the salts. Two scripts were utilized, linking several functions available within CERIUS2 in an automated fashion. Both used the molecular mechanics Drieding 2.21 MM force field, but varied in the method employed (Gasteiger or QEq) used for the assignment of atomic point-charges.19,20

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ized and compared in Table 5. It is clear from these data that the 25 counter-ions gave various outcomes in the salt screen. The following general categories of behavior were observed: (1) Strong acids (pKa < 2) gave salts from both water and methanol: nine strong acids (inorganic or sulfonic) gave salts from water, and the six of these that were evaluated in methanol gave the same salts in this solvent. This outcome follows the expectation discussed above concerning speciation. (2) Weak acids (pKa > 3), which in this study were carboxylic acids, gave salts from water but not from methanol. As discussed above, this is the result expected on the basis of the pKa shift of carboxylic acids in methanol from 3–5 to 8–10, so close to the pKa of ephedrine (9.74) that the solution chemistry precludes salt formation. (3) Some experiments yielded temporarily amorphous materials: malonic, succinic, glycolic, and L-malic acids all gave amorphous residues from methanol that crystallized over a period of time (1 month). Of these, crystal structures were obtained for all but the succinate salt (a monohydrate) which was disordered and for which, no satisfactory structure solution could be found. (4) Some systems yielded persistently amorphous material: formic, salicylic, trichloracetic, and fumaric acids gave amorphous solids from methanol; benzoic, citric, naphthalene sulfonic, and gluconic acids gave amorphous solids from water. In none of these cases was subsequent crystallization observed. (5) Some acids did not produce salts at all: in the cases of formic, salicylic, trichloroacetic, and fumaric in water this was thought to be related to insolubility, however, even dissolution in methanol did not enable salt formation.

Crystallization of pure ephedrine from water, ethanol, and methanol without special drying precautions gave highly crystalline material with a unit cell consistent with the hemihydrate previously reported.9 Hemihydrate crystals were dried over silica gel and re-analyzed, revealing a new, anhydrous form for which the structure was also determined.16 The overall results of the salt crystallizations from water are given in Table 3, and from methanol in Table 4. The outcome of this series of experiments proved repeatable in the 96-well plate. The results from both solvents are summar-

The three hydrates, maleate, succinate, tartrate did not give crystalline anhydrous forms. It may be that salt formation in these cases is only possible in the presence of water. It should be noted that it is extremely difficult to exclude water completely in small-scale evaporative crystallizations from hygroscopic solvents such as methanol. Although a systematic polymorph screen for each

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Table 3. Results Summary for Aqueous Crystallizations Crystal Growth Method/Outcome

pKa in Water

pH

4.76 4.19

6.05 5.89

Evaporation Solid glass

Formic Salicylic Trichloroacetic Adipic Fumaric Maleic

3.75 2.97, 13.82 0.66 4.44, 5.44 3.03, 4.38 1.92, 6.23

— — — 5.21 — 3.48

Nonsoluble Nonsoluble Nonsoluble Evaporation Non-soluble Evaporation

Malonic

2.83, 5.70

7.63

Succinic

4.19, 5.64

7.55

3.13, 4.76, 6.39 3.76

8.09 5.43

Crystallized from amorphous phase Crystallized from amorphous phase Soft gel Soft gel

3.82

5.30

L-(
)-malic

3.46, 5.10

3.94

L-(þ)-tartaric

3.02, 4.36

3.38

Acid Acetic Benzoic

Citric D-(
)-gluconic

Glycolic

Crystallized from amorphous phase Crystallized from amorphous phase Evaporation

7.09

Hydrochloric Nitric Phosphoric Sulfuric Benzene sulfonic Ethane-1,2-disulfonic Ethane sulfonic Methane sulfonic Naphthalene-2-sulfonic p-Toluene sulfonic


6.1
1.32 1.96, 7.12, 12.32

3.20 6.88 4.06

From solution From solution From solution


3.0, 1.92 0.70
2.06,
1.50
2.05
1.20 0.17

0.18 2.56 0.93 1.09 1.12 3.10

From solution From solution From solution From solution From solution Solid glass


1.34

1.31

From solution

Salt Name * Disordered Structure * Ephedrine acetate Ephedrine benzoate (amorphous) — — — Ephedrine adipate — Ephedrine maleate monohydrate Ephedrine malonate * Ephedrine succinate hydrate * Ephedrine citrate Amorphous Ephedrine D-gluconate amorphous Ephedrine glycolate Ephedrine L-malate Ephedrine L-tartrate monohydrate Ephedrine L-tartrate trihydrate * Ephedrine L-tartrate hydrate * Ephedrine hydrochloride Ephedrine nitrate Ephedrine dihydrogen phosphate Ephedrine bisulfate Ephedrine besylate Ephedrine edisylate Ephedrine esylate Ephedrine mesylate Ephedrine napsylate Amorphous Ephedrine p-tosylate

salt was beyond the scope of this study, no polymorphs were observed. EphedrineHCl salt, for example, was prepared from a range of solvents (DMF, DMSO, acetone, methanol, water, ethanol, 1-propanol, 2-propanol, 1-butanol, 1-pentanol, 1-hexanol, 1-octanol). All products had a needle (b-axis) morphology except methanol (plates) and water (variable-blocks, needles, and plates) and all PXRD patterns were consistent with the known ephedrineHCl salt crystal structure.10 Overall, these results highlight the sensitivity of salt screens to crystallization conditions. In

particular, they demonstrate the importance of—counter-ion solubility, effect of solvent on pKa, stoichiometry (for diacids), the timing of the end of the experiment, and the presence of water.

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PROPERTIES CORRELATIONS Physical Properties The physical property data obtained in this study are summarized in Table 6. Certain salts do not have thermal data in this set. For example, TGA

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Table 4. The Results Summary for Methanol Crystallizations Acid Acetic Benzoic Salicylic Fumaric Malonic Succinic Glycolic L-(
)-malic L-(þ)-tartaric Hydrochloric

Nitric Phosphoric Benzene sulfonic Ethane sulfonic Methane sulfonic

Experimental Outcome

No crystallization was observed in solution. Complete evaporation of all solvent resulted in a waxy precipitate of unreacted acid and base. This was confirmed with DSC and FTIR

Some crystalline material was observed in solution. Evaporation of all solvent resulted in flat, crystalline plates of the salt, confirmed with DSC, FTIR and XRPD

Some crystalline plates and needles were observed in solution. Evaporation of all solvent resulted in further crystalline plates and needles of the salt confirmed with DSC, FTIR, and later with XRPD (via pattern matching with water experiments)

showed that all four hydrates lost water prior to melting, so no melting data are included for these compounds. Likewise, DSC showed decomposition close to the melting point for the sulfate and edisylate salts, and multiple melting events for the nitrate and esylate salts, so no melting data were included for these four salts either. Correlations were then sought between the measured and calculated parameters. It was of particular interest to explore those relationships, which might be of practical use. For example, correlations between calculated properties such as lattice energy and density, which are readily available from single crystal data, and measured parameters such that solubility, heat of melting, and melting point are discussed below.

an un-buffered system) thus preventing the crystallization of the free base. As expected, the carboxylic acid salts have a less dramatic effect on pH, but interestingly, they tend to give higher solubilities than the salts of stronger acids. The measured solubilities (M/L), together with the solubility of the free base, are also shown in Figure 3a as a function of final, saturated solution pH. The solubility of the free base (the average of anhydrous and hemihydrate) has been calculated, as a function of pH, using the formula: ½B þ ½BHþ  ¼ Ksp ð1 þ ð½Hþ =ð10
pKa ÞÞ where [B] þ [BHþ] is the total concentration of ephedrine (charged and uncharged) in solution.1 At pH >> pKa, this simplifies to: ½B ¼ Ksp

Solubility and Final pH The solubility data (Tab. 6) show a wide range of values, from 0.06 to 1.65 g/L. It is clear that one benefit of acidic counter-ions is to lower the pH (in

Hence Ksp can be obtained from the solubility data for ephedrine free-base in Table 6. It is obvious from Figure 3a that the measured salt

Table 5. A Summary of all Crystallization Results Acids

Water Results

Methanol Results

DpKa Water

DpKa Methanol

Strong Inorganic Sulfonic Weak Carboxylic Dicarboxylic Hydroxy

9/10 crystalline

6/6 crystalline

8 to 12 units

2 to 6 units

8/15 crystalline including: 3 hydrates 4 conversions from an initial amorphous phase

0/9 crystalline

4 to 6 units


2 to 1 units

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Table 6. The Measured and Calculated Physical Properties of Ephedrine Salts

Abb. fb fbh ac ad mm mln gly lm ltm ltt hcl nit dhp bis bes eds es mes tos

Salt Free base Hemihydrate Acetate Adipate Maleate monohydrate Malonate (2:1) Glycolate L-malate L-tartrate monohydrate L-tartrate trihydrate(2:1) HCl Nitrate DHP Bisulfate Besylate Edisylate(2:1) Esylate Mesylate Tosylate

MW

Tm (8C)

DHm (kcal/M)

Density (g/cm3)

Elatt (kcal/M)

Sol (M/L)a

pH

165.23 174.24 225.28 311.37 299.32 434.52 241.28 299.32 333.33 534.68 201.69 228.25 263.22 263.31 323.4 520.65 275.36 261.33 337.42

34.7 * 100.4 105.5 * 182.6 105.4 125.1 * * 217.9

2.86 * 7.74 7.24 * 10.24 7.59 9.07 * * 8.7

1.114 1.157 1.191 1.207 1.273 1.235 1.302 1.31 1.349 1.279 1.229 1.285 1.394 1.44 1.316 1.335 1.237 1.329 1.294


20.57
25.13
40.94
68.43
69.61
40.43
61.91
62.03
104.46
94.06
57.58
54.02
66.52
69.77
62.72
43.15
47.26
59.25
61.12

0.345 0.643 3.751 5.299 3.792 3.392 1.828 1.774 1.761 3.146 1.601 5.770 1.417 0.896 0.260 1.152 1.493 2.246 0.323

11.52 11.34 5.98 5.4 3.41 7.85 5.22 3.74 3.38 7.1 3.18 6.85 4.06 0.17 2.91 0.86 1.28 1.19 1.33

b

b

180.3

7.1

c

c

184.3

10.26

c

c

b

b

143.3 156.7

8.76 9.29

‘Density’ refers to the value calculated from the crystal structure determination. *Hydrate. a Moles ephedrine per liters. b Multiple melting events. c Decomposes before melting.

solubilities all lie to the left of the free base solubility curve, indicating that all of the salt solutions are stable with respect to precipitation of the free base.

Here, we extend this concept to describe species that dissociate in solution, in which case, for complete dissociation we can equate the ideal solubility to the solubility product.22 Data were available for ephedrine-free base, nine 1:1 salts, and one 2:1 salt. Two of the 1:1 salts, adipate, and L-malate, are incompletely dissociated at the pH of solubility determination. For the ephedrine and the other eight salts, the ideal solubilities were calculated in the following way. The enthalpy of fusion per mole of ions was calculated by dividing the molar enthalpy of fusion in Table 6 by the

number of ions per molecule (1 for ephedrine, 2 for 1:1 salts, and 3 for the malonate 2:1 salt). This value along with the melting temperature was substituted into the van’t Hoff equation to give the ideal solubility expressed as a mole fraction. This quantity was converted to solubility in moles per liter for comparison with the measured data, as shown in Figure 3b. The ratio, g [(ideal solubility)/ (measured solubility)] was calculated as a guide to deviations from ideality. The data for ephedrine free-base are not included in the graph but its ideal solubility is 47.7 mol/L, 137 times the measured solubility of 0.34 mol/L. Thus, despite its low melting point, neutral ephedrine is relatively insoluble in aqueous solutions, with a strong positive deviation from ideality. This is presumably in large part due to the difficulty of solvating the phenyl ring in water. All the salts show much smaller deviations from ideality than the free base, and only one (the 2:1 malonate salt) shows a negative deviation (g ¼ 0.64), presumably because of the extensive solvation of the doubly charged anion. The two salts that show the strongest positive deviations from ideality are the besylate (g ¼ 15.6) and tosylate (g ¼ 10.5). These are the only two salts

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Solubility, Melting Point, Heat of Melting, and Lattice Energy The expected relationship between melting temperatures, Tm, melting enthalpies, ~Hm and ideal solubilities, xideal is described by the van’t Hoff equation:21 lnxideal ¼
DHf ðTm
TÞ=RTm T

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Figure 3. Various correlations for ephedrine solubility data. (a) The relationship between solubility (moles ephedrine per liter) and pH for ephedrine salts. The continuous line shows the solubility of the free base as a function of pH. Labels defined in Table 6. (b) The measured versus ideal solubilities of ephedrine salts. Labels defined in Table 6. (c) The variation of solubility (ln Ksp) with melting temperature for 9 1:1 ephedrine salts. Labels defined in Table 6. (d) The variation of solubility (ln Ksp) with enthalpy of fusion for 9, 1:1 ephedrine salts. Labels defined in Table 6. (e) The variation of solubility (ln Ksp) with lattice energy for ephedrine free base (fb), hemi-hydrate (fbh), and 17 salts.

with anions containing phenyl rings. The besylate salt is actually less soluble than the free base suggesting that the solubilizing effect of forming an ephedrine cation is overwhelmed by a combination of the higher melting temperature and enthalpy of fusion of the salt, combined with the relatively weak solubilization of the besyate anion.

The 1:1 salt that is closest to ideal solubility is the HCl salt (g ¼ 1.9). These results suggests that in general, for the ephedrinium ion in aqueous solutions the solubilizing effect of a single positive charge does not entirely overcome the counteracting difficulty of hydration of the nonpolar residues.

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Figure 3c and d shows the relationships between the solubility product, Ksp and the melting temperature and heat of melting respectively, for the nine 1:1 salts. To avoid using inconsistent units of Ksp the 2:1 malonate salt and the free base are omitted from these data. The expected qualitative correlations between solubility (ln Ksp) and melting data were observed, except for the phosphate and hydrochloride salt (see Section 5.4), with solubilities tending to decrease with increasing melting temperature and enthalpy of fusion. The quantitative relationships reported in previous studies were not found here. For pharmaceutical applications, salts are often investigated with the dual aims of raising the melting temperature and increasing the aqueous solubility. These data suggest that these aims are likely to be contradictory. The practical utility of determining the crystal structures for all the salts might be assessed by testing the relationship between the calculated lattice energy (a direct consequence of the crystal structure) and the measured solubility. This data set, Figure 3e, shows no correlation whatsoever. Overall, it appears that the best guide to solubility is the intuitive idea that the counter-ions, which are themselves strongly solvated, give corresponding salts with the highest solubilities. Given the weak correlations with heat and temperature of melting, it would seem that indeed solution chemistry, rather than crystal chemistry, is the determining factor in such a sequence of salts. Enthalpy of Fusion and Temperature of Melting The melting points and molar heats of fusion of this dataset are plotted in Figure 4. For the

2:1 malonate salt, the heat of fusion is expressed per mole of ephedrine. Data from the previous study of diastereoisomers6,7 are also included. The data show qualitative agreement with the hypothesis that higher melting temperatures correlate with higher enthalpies of fusion. Eight of the data points from this study show a reasonable correlation (R2 ¼ 0.92). The two outliers are both salts of strong inorganic acids; the HCl salt (Tm ¼ 4958) and the dihydrogen phosphate salt (Tm ¼ 4538C)—see also Figure 3c. The data point for ephedrine free base (Tm ¼ 307.85K, DHm ¼ 11.98 kJ/mol) is not included in the graph, but falls below the line of best fit. The data from the diastereoisomer study,6,7 which contained exclusively anions of aromatic derivatives of chiral cyclophosphoric acids, are not consistent with this quantitative trend. Interestingly (see also Fig. 4), they fall on a separate line which also includes the data points for the phosphate and HCl salts prepared here. These nearly linear relationships suggest two distinct values for the melting entropies of these two sets of salts. This is contrary to the outcome of Westwell et al., who showed a very poor correlation between melting enthalpy and melting temperature.23 Their data, however covered a much wider range of melting temperatures and incorporated both molecular and inorganic solids, but no molecular salts. Lattice Energy and Density Figure 5 shows the expected increase of calculated lattice energy with crystal density. The two largest lattice energies, however correspond to the tartaric acid monohydrate and trihydrate, which have lower densities than expected from the general trend. These are also the two lattice energies with the largest contribution from hydrogen bonding, which as in the case of ice, may be the reason for their relatively low densities. Otherwise, the trend of higher lattice energies at higher densities is as expected, but of no practical use. Lattice Energy and Enthalpy of Fusion

Figure 4. The enthalpy of fusion versus melting point for ephedrine salts: Filled squares—this work (HCl and DHP salts labeled). Open triangles—data from Leusen et al.6,7

Figure 6 shows the lack of correlation between lattice energy and enthalpy of fusion. It also emphasizes that the lattice energy is 5–10 times greater than the enthalpy of fusion, presumably due to contributions from ionization and vaporization enthalpies necessary to convert charged species in the melt into neutral species in the

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Figure 5. The variation of calculated lattice energy with crystal density for 19 crystal structures of ephedrine compounds. The three hydrates are labeled—see Table 6 for key.

vapor.23 Thus it can be concluded that this direct reflection of the precise crystal structures, the lattice energy, does not correlate with the measured parameter, the heat of fusion. This reinforces the conclusion drawn above for solubilities in which it was shown that there was no correlation between lattice energy and the solubility data. It is finally worth noting that, in contrast to previous work,6 we found the data points for structures with extended and folded molecular conformations to be randomly scattered within the correlations that we explored.

Comparison with Previous Property Correlations for Molecular Salts

Figure 6. The variation of calculated lattice energy with experimental enthalpy of fusion for ephedrine free base (fb) and 9 ephedrine salts.

Gould claimed (his Fig. 3), a linear relationship between log10 solubility and the reciprocal of the absolute melting temperature, with two of the six data pairs excluded.4 The pH of the solubility measurements was not given. This is not consistent with our current data, which indicate only a qualitative link between higher melting points and lower solubilities. In a study of 13 salicylate and chloride salts of chiral and racemic ephedrine and related compounds, Li et al.14 reported correlations between melting temperature and Ecoul (R2 ¼ 0.78) and also between enthalpies of fusion and EvdeW (R2 ¼ 0.75). The corresponding R2 values for the data obtained here with a larger range of counter-ions are R2 ¼ 0.20 and R2 ¼ 0.10 respectively (not shown here but see reference [15]). In summary, none of the correlations proposed by these previous authors was confirmed for this larger data set. The other correlation reported previously6 is between lattice energy differences and solubility ratios of diastereoisomers. It seems reasonable that enthalpies of ionization and vaporization, as well as deviations from ideality, will be similar for diasteroisomer pairs. However, these are not reasonable assumptions for salts of a common cation—the enthalpies of ionization and vaporization of the various counter-ions will vary and would need to be measured or estimated in

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order to account for them. This lies outside the scope of this present study.

DISCUSSION Salts, Solubility, and Salt Selection In many applications, including in vivo, the pH is controlled externally and, as seen in the forgoing discussion this will affect the salt solubility. Of particular interest may be to explore how the solubility changes between acidic and neutral environments. The behavior of ephedrine and selected salts in such circumstances can be discussed with reference to Figure 7. This shows the solubility of the HCl, malate, and besylate salts as functions of pH, calculated as described previously.18 At pH < 9, ephedrine free base will not precipitate unless the solution concentration is extremely high. If HCl is present, the solution concentration of ephedrine at pH 1–7 is determined by the solubility of the HCl salt. The behavior of the besylate salt demonstrates that the solubility of a salt can be lower than that of the free base—the besylate salt is less soluble than the free base at all pHs < 10. The fate of the L-malate salt is more complicated because the anion can exist as both single and doubly deprotonated species. At pH < 4, the solubility rises with falling pH as the dissolving monoprotonated acid is further protonated. At pH > 5, the solubility also rises as the dissolving monoprotonated form is deprotonated to the free acid. The minimum solubility occurs at pH 4.4, when the proportion of monoprotonated form is maximized and the solubility limited by the solubility product. The overall result of these complications is that on

passing from low to neutral pH, there is the possibility of precipitation and subsequent redissolution of the L-malate salt. From a product perspective, this may be highly undesirable since such uncontrolled physical changes occurring in vivo may destroy a carefully designed formulation and adversely affect the uptake of the drug into the blood stream. This plot highlights the significance of the pKa of the base. For a weak base with pKa  7, solubility will be very sensitive to pH at physiological pH values. For very weak bases (pKa < 5) with low solubility, the solubility at pHs in the range 5–7 will be determined by the solubility of the free base and will not be altered by salt formation. Salt Preparation—Importance of the Ternary Phase Diagram Difficulties are frequently encountered when scaling-up ‘‘hits’’ from a salt screen. A specific challenge in salt formation is to avoid unwanted precipitation of the free acid or the free base. A combination of the data on free base and salt solubility determined here, with literature data on the solubility of the acids24 has relevant implications. In such cases, plots of solubility as a function of pH can be misleading. A more rigorous alternative approach is to treat the system explicitly as containing three components, and use isothermal triangular phase diagrams. This methodology is commonplace for solventenantiomers-racemate systems,25,26 and is developed here for the analogous problem of salt formation. To aid with interpretation of the three diagrams presented here, it is recalled that:

Figure 7. The variation of solubility (total ephedrine in solution) with pH for 3 salts of ephedrine.

(1) The relative concentration of any component is 100% at the corner labeled with that component, and decreases linearly to 0% on moving away form that corner to the opposite edge of the triangle. (2) The solid phases all have fixed composition, whereas the liquid phase (denoted L) has variable composition. (3) The phase diagrams map out areas of compositions in which different solid and liquid phases will co-exist at equilibrium (4) In order to isolate the solid salt (S) robustly and consistently, the composition of the system should be such that the salt is in equilibrium with the liquid and no other solids.

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Table 7. Aqueous Solubility data (mol/L) at 258C

Acetic L-Malic Adipic

Acid

Salt

Ephedrine

— 10.4 0.2

3.8 1.8 5.3

0.5 0.5 0.5

This approach is illustrated for three ephedrine salts, the acetate, L-malate, and adipate. Solubility data in water for ephedrine free base and the three salts (from this work) and the corresponding acids22 are given in Table 7, and the schematic phase diagrams are shown in Figures 8a–c. Figure 8a shows the ternary diagram for the acetic acid/ephedrine/water system. The intersection along the ephedrine-water axis corresponds to the solubility of ephedrine in water. Acetic acid is a liquid at 258C and is completely miscible with water and hence there is no such intersection on the water acetic acid axis. The stoichiometry of the salt is 1:1 which defines the intersection on the ephedrine/acetic acid axis. The ternary diagram is then divided into various regions
L in which all compositions correspond to single phase solutions; E þ L where solution is in equilibrium with crystalline ephedrine; S þ L in which the crystalline salt is equilibrium with solution; S þ E þ L where both solid phases can co-exist with solutions. This information greatly simplifies the definition of a process for making the acetate salt. In the presence of excess base, there is a risk that free base will crystallize. This could occur inadvertently as a consequence of evaporation of acetic acid during crystallization or due to addition of a solution containing acid to one containing the base. However, in the presence of excess acid, only the salt can crystallize (‘‘S þ L’’

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region). Hence, it is preferable to maintain an excess of acetic acid when preparing ephedrine acetate. Figure 8b shows the phase diagram for the water-ephedrine-L-malic acid system. L-malic acid is a solid at room temperature and this introduces two additional regions into the ternary diagram
M þ L and M þ S þ L in which the crystalline free acid (M) can appear. The high solubility of malic acid means that once again it is relatively easy to crystallize only the salt provided that excess acid is present (‘‘SþL’’ region). Note that careful washing of product crystals will be necessary to avoid the mother liquors entering the ‘‘M þ S þ L’’ region during drying, with concomitant formation of some solid malic acid. Finally, Figure 8c shows the schematic phase diagram for the ternary system water– ephedrine–adipic acid. In this case, the preparation of ephedrine adipate is made much more complicated by the low solubility of both ephedrine and adipic acid which compresses the ‘‘S þ L’’ region. Tight control of stoichiometry is required to crystallize only the salt (‘‘S þ L’’ region) and to avoid crystallization of either the free acid (‘‘A þ S þ L’’ region) or the free base (‘‘S þ E þ L’’ region). Moreover, careful mixing of the reagents is necessary, and becomes more crucial on scale-up where longer addition times increase the risk of precipitation of unwanted solids, which may not redissolve. The comments above about careful washing also apply here. These schematics phase diagrams are relevant to salt synthesis from water at pHs where speciation is appropriate to salt formation. Where an organic solvent is used, the corresponding solubilities of base and acid in that solvent would need to be considered. The common practice of

Figure 8. Schematic ternary phase diagrams for the systems (a) ephedrine (E)–water–acetic acid, (b) ephedrine (E)–water–malic acid (M), (c) ephedrine (E)–water–adipic acid (A), including the salt (S). DOI 10.1002/jps

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using different solvents for acid and base is obviously more difficult to analyze. It is interesting to note that despite the numerous reports of salt screens and screening protocols, we could find no example of a measured ternary phase diagram for this type of system, nor an appreciation of its importance. In addition, it has not escaped our attention that the ternary diagrams described here for salt preparation will also be relevant in the currently active field of co-crystals. Consequences for Salt Screening Protocols In some salt screening protocols, stock solutions of acids in alcoholic solvents are used to facilitate automated dosing. These are preferred to aqueous solvents because addition of water often causes the free base to precipitate. This study has demonstrated that while this approach is acceptable for strong acids it is flawed for weak acids such as carboxylic acids. Salts, particularly of weak acids, that form readily in water can be missed completely in nonaqueous solvents. Salt screening is often carried out precisely because the free base has low solubility in water. In particular, it may not be possible to make up a ‘‘stock’’ aqueous solution of sufficient strength, as required for many automated procedures. The data presented here suggest that it would be advantageous to add solid free-base to an aqueous solution of the acid, or vice-versa, in order to achieve solubilization and subsequent reaction of acid and base species to form a salt. This illustrates one potential disadvantage of allowing automation requirements to dictate screening protocols. A further observation of our current studies was that in four cases out of 25, salt crystallization from the initial waxy precipitate did occur but over very long time periods. This evolution to the crystalline state would have been missed by standard salt screening protocols, in which the samples are generally analyzed at only one time point, usually less than 1 week after evaporation is complete.

and to a lesser extent racemates and diastereoisomers, where there have been successes in correlations of this type.27,28 In these cases, many assumptions can be made which do not apply to salts of a common counterion. In particular, for a series of salts deviations from ideality will not be constant, and ionization and vaporization energies must be considered when relating lattice energy calculations to solid–liquid equilibria.

CONCLUSIONS

Overall, we have to conclude that, in fact we find no correlation between the detailed crystal structure data, as reflected in the lattice energies, and any measured parameter. This outcome implies that while full structural characterization of salts is essential, the exclusive use of structural information as a guide to salt selection has little to recommend it. This is in contrast to polymorphs

This study, detailing the outcomes for a salt screen of 25 acids with the single base ephedrine, has resulted in significant new insights and clarified a number of highly important fundamental and practical issues surrounding this extremely common industrial practice. It is clear that the design of screening protocols should take into account solvent effects on solution speciation, should focus more on use of water as a crystallization medium, even if the active is insoluble and that time should be considered a key variable especially when waxy solids form on solvent evaporation. The combined measurement of solubility and solution pH has enabled the calculation of the relationship between these parameters across the entire range of ephedrine salts. This clarifies significant issues concerning salt selection and in particular, allows those salts which may have unfavorable pH-solubility profiles to be identified. Structure–property relationships, as reflected in such parameters as lattice energy, heat and temperature of melting, crystal density and solubility, follow the expected trends. A direct structure–property correlation between lattice energy and solubility, however, was not found. It appears that in this context and in aqueous solution, it is the process of ion solvation rather than the strength of solid state interactions that governs solubility. With hindsight, the pKa and the solubility of the counter-ion are much more useful guides than crystal structures and lattice energy calculations. Thus, the idea that high solubility might be associated with, say, a low melting point while roughly true, is not such a good guide as consideration of the potential for the solvent to solvate the chosen anion. Because it is not possible to describe the solubility of a salt as a function of pH using a single parameter, we now suggest that it is not wise to seek such linear correlations. This leads ultimately to the conclusion that in the

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design of solid phases, the application of crystal engineering approaches will be extremely limited unless supported by an equally robust appreciation of factors determining the solution phase interactions.

ACKNOWLEDGMENTS The authors would like to acknowledge the help given in numerous discussions with Brian Cox. Helen P Jones (AstraZeneca) and Julian Cherryman (Intertek) on aspects of Ksp and lattice energy calculation. The financial assistance of AstraZeneca, which funded the Ph.D. studentship (EC) and thesis on which this paper is based, is acknowledged.

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27. Bernstein J. 2002. Polymorphism in Molecular Crystals. Oxford: OUP. 28. Leusen FJJ, Noordik JH, Karfunkel HR. 1993. Racemate resolution via crystallization of diastereomeric salts—thermodynamic considerations and molecular mechanics calculations. Tetrahedron 49:5377–5396.

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