Systematic Data Set for Structure−Property Investigations: Solubility

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Systematic Data Set for Structure-Property Investigations: Solubility and Solid-State Structure of Alkaline Earth Metal Salts of Benzoates Jean-Baptiste Arlin,† Alastair J. Florence,‡ Andrea Johnston,‡ Alan R. Kennedy,*,† Gary J. Miller,† and Kirsty Patterson† † ‡

WestCHEM, Department of Pure & Applied Chemistry, University of Strathclyde, Glasgow G1 1XL, Scotland Solid-State Research Group, Strathclyde Institute of Pharmacy and Biomedical Sciences, University of Strathclyde, Glasgow G4 0NR, Scotland

bS Supporting Information ABSTRACT: A new resource for studying structure property relationships is presented, namely a systematic database of 36 organic salt structures together with phase specific aqueous solubility data. The salts are derived from four M2þ cations (Mg2þ, Ca2þ, Sr2þ, Ba2þ) and nine substituted benzoate anions. The intrinsic solubility of the free acid is found to have a major contribution to make to salt solubility, but despite previous literature assertions, there appears to be little correlation of solubility with the polarity of the organic ions, with cation size, or with hydration state. Importantly, we also show that consideration of the array structure rather than just molecular considerations improves prediction of rank orders of solubility. Thus, three-dimensional intermolecular networks (here formed with hydrogen bonding, M-O-M and M-N-M interactions, and halide interactions) are found to have lower aqueous solubilities than lower dimensional networks.

’ INTRODUCTION Salt selection in the pharmaceutical industry is the process of choosing a pharmaceutically acceptable counterion for a biologically active ion, to optimize one or more performance critical physicochemical properties. Solid dosage forms are preferred, and so in an ideal world, the salt selection process should be an exercise in rational solid-state chemistry—manipulating the different intermolecular interactions made by the different ion pairs to obtain the required solubility, dissolution rate, melting point, particle shape, or other such desired property.1 Unfortunately, the current understanding of the correlation between solid-state structure and physicochemical properties is inadequate for most practical purposes. Thus, salt selection is still carried out by inefficient and time-consuming trial and error methods involving the preparation and testing of numerous salt forms. A prime reason for this lack of understanding is that solid-state structure-property correlation analysis is hampered by the absence of suitable, large, and systematically connected structural data sets complete with comprehensive listings of phase specific physical properties. Until recently, single crystal diffraction measurements were relatively slow and tended to be reserved for compounds of some “special interest”. Only within the past decade has instrument capacity increased to a stage where it is practically possible to generate large numbers of crystal structures, to examine one particular problem, within reasonable time r 2011 American Chemical Society

scales.2 In the past, traditional solubility studies of differing salt forms were typically carried out with little or no knowledge of the solid-state structures of the salts being examined, and so physicochemical properties tended to be linked to molecular structure or “chemical identity” alone. Examples of such studies relevant to this work are the observation that hydrated forms of organic crystals are less water-soluble than their equivalent anhydrous forms,3 that aqueous solubility in salts (including specifically those of para-aminosalicylate) decreases with increased hydration state,4 and that, for s-block metal salts of certain carboxylated drugs, aqueous solubility increases with increase in charge and decreases with increase in ionic size.5 A problem with such studies is that they attempt to describe macroscopic properties based on molecular theory and in the absence of any information about array structure. Thus, general observations from one material are often apparently countered by contradictory behavior observed in a second material. For example, extra polar groups may increase water solubility,6 in line with the simple idea of polar entities being more soluble in the polar aqueous environment, but in other cases, addition of polar groups is observed to decrease solubility.7 The explanation for these apparent contradictions often lies in the “missing” structural detail. In the case above, more polar entities can favor Received: November 22, 2010 Revised: January 25, 2011 Published: February 18, 2011 1318

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Scheme 1. Molecular Structures of the Nine Benzoate Anions Used Throughouta

a

These nine anions are derived from benzoic acid (ben), o-aminobenzoic acid (onh2), salicylic acid (sal), p-fluorobenzoic acid (pf), p-chlorobenzoic acid (pcl), o-nitrobenzoic acid (onit), p-nitrobenzoic acid (pnit), p-aminobenzoic acid (pnh2), and p-aminosalicylic acid (pams).

intermolecular bonding in the solid-state as well as in solution.6 Anderson and Flora summarize the predictive problem by pointing out that chemical changes that increase free energies of ion hydration (e.g., increasing ionic charge, decreasing ionic radii, adding polar groups) also tend to increase lattice energy. “Thus the overall effect...on water solubility will depend on which terms...are most sensitive to the change.”5a Here we take the first steps toward alleviating the lack of structural understanding by presenting a set of data (36

systematically related solid-state structures with phase specific aqueous solubility data) which we believe will be of seminal use to those interested in examining the structure-property relationship between pharmaceutically relevant salt structures and solubility. Some other large systematic structural studies with different salt forms of drugs have been reported8 or could be assembled from the structural databases,9 but these typically lack the necessary phase specific physicochemical data needed for structure-property correlation. Such joint structural-property 1319

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Table 1. Coordination Network and Hydration State

Number of water molecules per metal atom. Given as number of water ligands, number of free water molecules of crystallization. Here “chain” is used to describe a 1D coordination polymer with a single row of cations, and “ribbon” is used for 1D coordination polymers with a double row of cations (such as those seen in ladder structures). a

Figure 1. Extended structure of Mgpnit viewed down the b axis. Note the structure's solvent-separated ion-pair nature and the organicinorganic layering. Here, and in other color figures, dashed bonds are hydrogen-bonding contacts; metal = green; O = red; N = blue; C = black; H = gray.

studies that do exist tend to be compound specific and fairly small,10 though a few notable, larger studies are now available.7,11,12 With 36 structure-property pairs, our data set is designed to give an experimentally based insight into structural effects on solubility trends in group 2 metal salts of ortho- and para-substituted benzoic acids. Mg, Ca, Sr, and Ba dications were chosen, as they represent a systematically changing set whose bonding is based on well described fundamental properties (essentially ionic radius and its effect on charge density). Mg and Ca are of course also common counterions for acidic drugs.1 The anions chosen were simple benzoic acid derivatives (Scheme 1). Although some of these have intrinsic pharmaceutical interest (e.g., salicylic and para-aminosalicylic acids) or indeed broad materials science applications (benzoic acid), they were largely selected as model weakly acidic compounds combining a chemical fragment commonly found in many pharmaceuticals (e.g., 5-fluoroquinolines, aspirin, naproxens) with the

Figure 2. View of the molecular structure of Mgonit showing the coordination geometry about Mg.

requirement of easy access to a systematically changing series of ring substituents.13

’ EXPERIMENTAL SECTION General. Group 2 metal salts of ben, sal, and pam were obtained by reaction of a nearly saturated aqueous solution of the appropriate Na salt with a slight excess of an aqueous solution of MCl2. After stirring for approximately 1 h, the colorless precipitate was collected by filtration and then recrystallized from water. All other salts were prepared by reaction of aqueous solutions of the appropriate carboxylic acid with slight excesses of metal carbonates. Reactions were stirred for at least 2 h. If, at the end of this time, some solid was still apparent in the solution, the reaction medium was heated until ebullition of solvent. Finally, the solution was filtered and left to evaporate in order to produce crystals. 1320

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Figure 4. Part of the 1-D coordination polymer formed by Mgben. Metal to metal bridging is solely through carboxylate groups, and the polymer is supported by hydrogen bonding.

Figure 3. Extended structure of Capnit showing the Ca dimer unit created by bridging water molecules. Note also that two of the benzoate units form Ca-O bonds but that the remaining two benzoates do not bond to Ca.

Crystallography. Of the 36 structures discussed, data on 12 were obtained from cif files supplied by the CCDC.9,14-22 For the remaining 24 structures, samples for single crystal diffraction studies were either obtained directly from the solutions prepared above or from simple aqueous recrystallizations based on slow evaporation or controlled cooling. Measurements were recorded at low temperature with Nonius Kappa CCD and Oxford Diffraction diffractometers with Mo KR radiation (λ = 0.71073 Å), except for compound Mgben, where data were obtained at station 9.8 of the Daresbury Synchrotron Radiation Source.23 All structures were refined against F2 to convergence using the program SHELXL-97.24 Selected crystallographic data and refinement parameters are given in the ESI and full details deposited as cif files. These can be obtained free of charge from The Cambridge Crystallographic Data Centre via www.ccdc.cam.ac.uk/data_request/cif by quoting deposition numbers CCDC 779284 to 779307. Powder X-ray diffraction data was collected using a Bruker D5000 diffractometer operating in reflection mode and using Cu KR radiation (λ = 1.54056 Å). Data was collected at ambient temperature in the range 4.000° to 35.056° in 2θ. Dry samples for powder diffraction were ground with an agate mortar and pestle to produce a homogeneous polycrystalline powder. However, samples recovered from the slurry used in solubility measurements were presented “damp” and as recovered. This was an attempt to avoid potential problems with phase changes on removal of solvent. Solubility Measurements. Samples for solubility measurements were repeatedly recrystallized before use. Chemical and phase purity prior to use were checked by microanalysis and powder diffraction (see the Supporting Information). Solid in excess was added to 10 mL of water to form a slurry. The slurry was left stirring at 25 °C in an incubator for approximately one week. Exceptions were the salts of pams, which were observed to decompose in solution over a period of days.4a Here the slurry time used was approximately 36 h. The pH of the saturated solution was determined using a pH meter (Hanna Instruments Piccolo 2). The slurry was initially filtered through a filter funnel using a Fisherbrand QL100 filter paper and then filtered through a syringe filter (Whatman Anotop 50, 0.2 μm pore size). The identity of the excess solid from the slurry was checked by powder diffraction (see the Supporting Information). In all cases only salt forms were recovered with no trace of free acids. To determine the solubility, the UV absorbance of aromatic compounds was used. Two calibration curves were established for each anion, within the linear domain of the Beer-Lambert law. An average of the two gave the linear relation linking the absorbance of the anion and

Figure 5. Part of the 1-D coordination polymer formed as a cationic entity within the structure of Capf. Anions not directly bound to Ca are omitted.

Figure 6. Packed structure of Capf viewed along the c axis and hence down the length of the 1-D metal coordination chains. These chains are linked by hydrogen bonding through the pf anions that are not bound to metal. Organic bilayers are formed. its molarity. Each calibration curve had to be composed of at least four points within the 0-1 absorbance domain and have a linearity index R2 > 0.999, or the measurements were rejected. By diluting the saturated solution (obtained above) until its absorbance was within the linear domain, the saturated solution molarity was calculated. Each salt’s solubility value was obtained at least in duplicate.

’ RESULTS AND DISCUSSION Formation of the Structural Data Set. Single crystal structures of all four metal salts (Mg, Ca, Sr, and Ba) were obtained for nine of sixteen originally selected benzoic acids; see Scheme 1.25 Only the 36 compounds from these complete quartets were considered further, of these, 24 structures are reported here for the first time, and the structures of the remaining 12 were 1321

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Figure 7. Part of the 1-D coordination polymer present in the structure of Capnh2.

available in the literature14-22 and were used as obtained from the Cambridge Structural Database.9 Where multiple phases of a given salt were identified, only the form that corresponds to that recovered from the slurry used to measure aqueous solubility is presented herein. Coordination Networks. The structures were first examined from the viewpoint of metal bonding, traditionally seen as being of prime importance in directing local coordination and packing in metal-organic chemistry. Table 1 shows the hydration states and metal coordination-polymer network-dimensionality and type. It can be seen that the 36 salts offer a wide range of structures, but inspection of the data reveals that there are relationships present—structural trends that are discussed below. Overall, none of the salt structures contain free acid molecules (as seen, for instance, in [K(aspirin)(Haspirin)]).26 With the exception of anhydrous Bapf, all are hydrates with at least one water molecule directly bound to the metal ion. The general formula is thus [MLx(OH2)y][L]2-x 3 zH2O (x = 0, 1, or 2; y = 0-7; z = 0-4), and the total hydration state ranges from 0 to 9. Mg salts tend to have the highest hydration state for each quartet of carboxylate salts (true for 6 out of 9 acids). This can be related to the magnesium ion’s small size and hence high charge density and strong oxygen-bonding nature,13a but there does not seem to be any indication that the reverse is true; that is, the larger Ba ion does not seem to form generally lower hydration state salts than Ca or Sr. The pnit quartet all have high hydration states that result in the Ca and Sr pnit salt structures being of a different structural type to those of the other Ca and Sr salts. Apart from this, there is little difference apparent in hydration behavior across the different benzoate derivatives. The Table 1 color scheme highlights some general similarities in structural types identified across the salt library. Key features of the five larger groups are given in more detail below, but general features are as follows. The groupings are largely metal ion dependent. Thus, only Mg forms solvent-separated ion-pair structures. With the exceptions of the anomalous Capnit and Srpnit structures noted above, it is also only Mg that forms discrete structures, where no coordination polymer based on M-O bonding occurs. The Ca and Sr salts tend to exhibit mutually similar structural types. This is emphasized by the presence of three near isomorphic structural pairs, those for pnh2, sal, and pams. Bapcl is approximately isostructural with Srpcl, but otherwise, the Ba structures are highly individual and largely defy grouping. Their salts do tend to form higher dimensional metal coordination networks than the lighter (and smaller) metal equivalents, though Bapams is an exception to even this simple observation. Metal Coordination Based Structural Classes. i. SolventSeparated Ion-Pairs. Three Mg salts form solvent-separated ion-pair structures with formula [Mg(OH2)6][L]2 3 2H2O (see Figure 1).

Figure 8. Packed structure of Capnh2 viewed along the length of the metal coordination chains. Note that the alternating organic-inorganic layering seen in other systems is not present here.

These are Mgonh2, Mgpnh2, and Mgpnit. All three structures feature alternating layers of organic and inorganic nature. The hexaaquamagnesium cation is a common counterion to organic anions;27 indeed, in salts of stronger acid types, it is often the only Mg species seen.13,28 Throughout this work, only the Mg ion forms complexes with no metal to OOCR bond. Previous work on sulfonic acids has suggested that Mg has a preference, not seen in the heavier metals, for bonding to the neutral water ligands over the formally charged acid group because it is the least electropositive and hence most covalent of the four group 2 metals used here.13 ii. Discrete Structures. Again three Mg structures (Mgonit, Mgsal, and Mgpams) fall into this class. The general formula is [MgL2(OH2)4] with mutually trans carboxylate ligands each bonding to octahedral Mg through one O atom (see Figure 2). In each case, despite the presence of ortho substituents capable of forming chelate interactions, there are no further Mg to anion interactions and so no coordination polymer forms. Alternating organic and inorganic layers are again observed, but now the organic layer is a bilayer. Two other salts, Capnit and Srpnit, also give discrete structures. However, these differ from those of the Mg group. Each of these structures feature one η2 chelated carboxylate anion and one carboxylate anion that does not bond to metal (see Figure 3). Together with high numbers of both water ligands and waters of crystallization, they have dimeric and monomeric structures, respectively [(H2O)4(pnit)Ca(μH2O)2Ca(pnit)(H2O)4][pnit]2 3 8H2O and [Sr(OH2)7(pnit)][pnit] 3 2H2O. A further difference with the discrete Mg structures is that neither forms an organic bilayer. iii. Mg Coordination Polymers. In contrast to the previously described structures, the remaining three Mg salt structures do form coordination polymers through bridging COO groups. Mgben and Mgpcl are similar and are both based on octahedrons with three fac H2O ligands and three bonds to O atoms from COO groups. One of these carboxylate bonds is to a terminal acid fragment, but the others form 1-D polymeric chains through Mg-O-C-O-Mg linkages (see Figure 4). The third structure, that of Mgpf, is different. All of the acid anions bridge between metal ions, with the two terminal H2O ligands trans to each other. This forms a 2-D coordination polymer through Mg-O-C-O-Mg bonding. As with group ii above, organic bilayers are formed. 1322

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Table 2. Observed Aqueous Solubilitiesa salt mgben caben

formula (from SXD characterization). [MgL2(OH2)3] 3 H2O [CaL(OH2)3][L]

solubility (esd) mol L-1

salt

formula (from SXD characterization).

solubility (esd) mol L-1

0.591(29)

srpnit

[SrL(OH2)7][L] 3 2H2O

0.045(2)

0.096(1)

bapnit

[BaL2(OH2)5]

0.054(2)

srben

[SrL2(OH2)]

0.160(3)

mgpnh2

[Mg(OH2)6][L]2 3 2H2O

0.252(26)

baben

[BaL2(OH2)2]

0.061(1)

capnh2

[CaL2(OH2)2]

0.226(12)

mgpf

[MgL2(OH2)2]

0.144(1)

srpnh2

[SrL2(OH2)2]

0.717(14)

capf

[CaL(OH2)3][L]

0.012(1)

bapnh2

[BaL2(OH2)2] 3 0.45H2O

0.659(10)

srpf

[SrL(OH2)3.87][L] 3 0.87H2O

0.035(1)

mgsal

[MgL2(OH2)4]

0.702(11)

bapf mgpcl

[BaL2] [MgL2(OH2)3] 3 H2O

0.035(1) 0.145(1)

casal srsal

[CaL2(OH2)2] [SrL2(OH2)2]

0.079(4) 0.144(6)

capcl

[CaL(OH2)3][L]

0.016(1)

basal

[BaL2(OH2)]

0.420(9)

srpcl

[SrL(OH2)4][L]

0.035(1)

mgonit

[MgL2(OH2)4]

0.838(21)

bapcl

[BaL(OH2)4][L]

0.034(1)

caonit

[CaL2(OH2)2]

0.434(5)

mgonh2

[Mg(OH2)6][L]2 3 2H2O

0.047(2)

sronit

[SrL2(OH2)4]

0.629(8)

caonh2

[CaL2(OH2)3]

0.062(1)

baonit

[BaL2(OH2)3]

0.837(22)

sronh2

[SrL2(OH2)2] 3 H2O

0.165(1)

mgpams

[MgL2(OH2)4]

0.200(28)

baonh2 mgpnit

[BaL2(OH2)] [Mg(OH2)6][L]2 3 2H2O

0.039(2) 0.059(5)

capams srpams

[CaL2(OH2)0.5] [SrL2(OH2)0.5]

0.582(11) 0.260(3)

capnit

[CaL(OH2)5][L] 3 4H2O

0.082(1)

bapams

[BaL2(OH2)3]

0.251(3)

As no absolute prediction of solubility values is attempted, all concentrations are given in mol L-1. This is sufficient to rank solubilities. As all species measured are ML2 based, the order of benzoate “drug” concentrations is identical. Entries in italics indicate that PXRD of solid recovered from the slurry did not match with that expected from the SXD determination. a

Table 3. Rank Order Aqueous Solubility by Carboxylate

Table 4. Rank Order Solubilities by Metala

carboxylate

literature ranking

found anion

found anion

anion

of free acid solubilitya

ranking rangeb

average rankingb

carboxylate anion

solubility order

sal

Mg > Ba > Sr > Ca

onit pnh2

1 2

1-2 1-4

1.5 2.5

onit pf

Mg g Ba > Sr > Ca Mg > Ba = Sr > Ca

onh2

3

4-9

6.75

pcl

Mg > Ba = Sr > Ca

ben

4

3-5

4.25

ben

Mg > Sr > Ca > Ba

sal

5

2-6

4.25

pnit

Ca > Mg > Ba > Sr

pams

6

1-5

3.25

pnh2

Sr > Ba > Mg > Ca

pnit

7

5-8

6.5

onh2

Sr > Ca > Mg > Ba

pf

8c

6-9

8

pams

Ca > Sr > Ba > Mg

pcl

9c

8-9

8

a

a

Entries in italics indicate that PXRD of solid recovered from the slurry did not match with that expected from the SXD determination.

iv. Ca, Sr, and Ba 1-D Coordination Ribbons. Four structures, Srpf and the Ca salts of ben, pf, and pcl, are based on eight coordinate metal centers with identical ligand behavior. All these structures have one COO group that does not bond to metal and one COO group that makes four individual O-M bonds, bridging three metal ions (see Figure 5). One of the three independent H2O ligands also bridges between metals with the other two terminal. This results in 1-D coordination polymers which are further linked by hydrogen bonding with the nonmetal binding acid group (see Figure 6). Two more salts, Srpcl and Bapcl, have similar structures—but with an extra bridging water ligand. In all these compounds, the coordination ribbon is arranged so that all the aromatic groups pendant upon it are coplanar. Alternating organic and inorganic layers are seen, and

in all cases organic bilayers, as shown in Figure 6, are a prominent packing feature. v. Ca and Sr 1-D Coordination Columns. As with the previous group, these five structures are based on eight coordinate metal centers. However, here all of the COO groups bond to metal. There are only two mutually cis H2O ligands, and neither of these bridges between metals. Capnh2 and Srpnh2 have identical structures with one carboxylate ligand acting as seen in group iv to make four individual bonds (and hence all the metal to metal bridges). The second carboxylate chelates to the metals in a terminal fashion (see Figure 7). Three further structures, Caonit, Casal, and Srsal, are variations on this theme. All feature ortho substituents that sterically block one carboxylate O atom per ligand from easily making two bonds to metals. Instead, here both COO ligands each make three individual bonds to metal atoms, and thus, both are involved in bridging and hence propagating the coordination polymer. All structures in group v have carboxylate ligands projecting in four directions from the core inorganic chain. The packed structure thus does not have

Ranked as per values for room temperature aqueous solubility in ref 34. 1 = highest solubility, 9 = lowest solubility. b For each of the four metal cations, the nine carboxylate salts were again ranked from 1 to 9. For each carboxylate anion, the table gives both the range of rankings found across the four metal series and the average of these. c Variously described as “not soluble” or “slightly soluble” at room temperature.

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Table 5. Intermolecular Bonding Networks and Solubility for Mg and Ca Salts dimensionality of

dimensionality of

dimensionality of

salt

metal coordination networka

total intermolecular networkb

intermolecular network after modificationc

Mgonit

0

2

2

HB

HB

Mgsal

0

2

2

HB

HB

0.702(11)

Mgben

1

2

2

M-Ocarb HB

HB

0.591(29)

Mgpnh2

0

3

3

HB

HB

HB

0.252(26)

Mgpams

0

3

2

HB

HB(N)

0.200(28)

Mgpcl Mgpf

1 2

3 3

3 3

M-Ocarb HB M-Ocarb HB

HB(N) NH 3 3 3 π HB M-Ocarb HB

XB XB

0.145(1) 0.144(1)

bond type in first dimensiond

bond type in second dimensiond

bond type in third dimensiond

solubility (mol L-1)e 0.838(21)

Mgpnit

0

3

3

HB

HB

OH 3 3 3 O2N

0.059(5)

Mgonh2

0

3

3

HB

HB

HB

0.047(2)

Capams

2

3

2

MLM

MLM

HB(N)

Caonit

1

2

2

M-Ocarb HB

Capnh2

1

3

1

M-Ocarb HB

OH 3 3 3 O2N HB(N)

Caben Capnit

1 0

2 3

2 3

M-Ocarb M-Oaq HB

HB HB

Casal

1

3

3

M-Ocarb HB

HB

Caonh2

1

2

2

M-Ocarb HB

XB

0.016(1)

XB

0.012(1)

Capcl

1

3

3

M-Ocarb M-Oaq

HB NH 3 3 3 π HB

Capf

1

3

3

M-Ocarb M-Oaq

HB

0.582(11) 0.434(5)

HB(N) OH 3 3 3 O2N HB

0.226(12) 0.096(1) 0.082(1) 0.079(4) 0.062(1)

a

Dimensionality of network created by metal to O or N bonds. b Dimensionality of network created by all polar noncovalent intermolecular interactions. c See text for description of modifications made. d Main bonding types seen to propagate network. In most cases the three “dimensions” correspond to the three crystallographic axes—but in some cases diagonals to these axes are used. Key: HB = hydrogen bond of OH-O type; HB(N) = hydrogen bond involving NH2; M-Ocarb = metal to carboxylate bridges; M-Oaq = metal to water bridges; MLM = coordination network propagated by metal ions being linked by bonding to more that one substituent of the benzoic acid; XB = halide to haloarene interactions (X = F or Cl). e The standard deviation in the last decimal place is given in parentheses.

alternating organic and inorganic layers—instead inorganic channels are formed (see Figure 8). This channel structure is very distinct from the structures seen for the other salts. Aqueous Solubility. The equilibrium solubilities of each salt were measured by slurrying samples in deionized water at 25 °C for 7 days to achieve equilibrium conditions, filtering the sample, and then diluting the filtrate and measuring concentration by UV/vis spectroscopy. Results are shown in Table 2 and further analyzed in Tables 3-5. Powder X-ray diffraction (PXRD) patterns of the solid samples both before slurrying and of the excess solid recovered after slurrying were checked against the appropriate SXD structure. For four of the 36 salts studied (Bapnit, Baonit, Srpf, and Mgonh2), PXRD indicated that a solution-mediated phase transformation had occurred. For Bapnit and Srpf, single crystal structures were subsequently successfully obtained for the phase recovered from solution—e.g. the phase for which the solubility has been measured. However, this was not possible for Baonit and Mgonh2, and so here the solubility data measured is not that for the structural phase reported. These two structures are thus excluded from any discussion relating intermolecular bonding to solubility. The most obvious result and, we would argue the most important, is a negative result. Namely that the data set shows that a number of factors often invoked as having a large impact on solubility seem to have no determining effect here. The undergraduate textbook approach, e.g. simple inspection of the organic anions for number or strength (or indeed position) of polar substituents,5,6 does not correlate with observed solubility. Neither can any one metal cation (of whatever size) be said to

be consistently more or less soluble than the others,4,5 nor can any simple relationship to hydration state be seen.2-4 All three of these factors derived from the chemical identity of the solute are commonly quoted in the literature in attempts to rationalize solubility, but here once given a relatively large and systematic set of data, they fail. Study of Table 3 shows that some relationship to the solubility of the free acid forms is present. Thus, Honit and Hpnh2 are the most soluble free acids and give salts with generally high apparent solubilities. Similarly, Hpf and Hpcl give low solubilities for both their free acid and salt forms. Various studies have previously suggested that the intrinsic solubility of the parent species is an important predictor of salt solubility.10c,11,29 That this is far from the whole story here can be seen in the detail. There are large discrepancies between free acid solubility rank order and the average salt rank order for pams and onh2. There are also many individual salt forms whose rank solubilities lie far from those implied by the free acid rank order. Moreover, while free acid solubility is obviously important, from the point of view of pharmaceutical salt selection, considering free acid solubilities of the organic portion tells us nothing about which salt form of a given acid will be most soluble. Solubility and Cation. Ranking salt solubility for each carboxylate gives the cation rank orders shown in Table 4. Again, the important feature here is a negative result. In this case, it is that none of the nine lists follow group 2 of the periodic table (e.g., neither order Mg > Ca > Sr > Ba nor Ba > Sr > Ca > Mg is present). This suggests that fundamental cation properties that follow periodic trends (ionic radii, electronegativity, charge 1324

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Table 6. Intermolecular Bonding Networks and Solubility for Sr and Ba Salts

salt

dimensionality of

dimensionality of

dimensionality of

metal coordination networka

total intermolecular networkb

intermolecular network after modificationc

bond type in first dimensiond

bond type in second dimensiond

bond type in third dimensiond HB(N)

solubility (mol L-1)e

Srpnh2

1

3

1

M-Ocarb HB

HB(N)

Sronit

1

2

2

M-Oaq

HB

0.717(14)

Srpams

2

3

2

MLM

MLM

Sronh2

1

2

2

M-Ocarb M-Oaq

HB

Srben

2

2

2

M-Ocarb M-Oaq

M-Ocarb HB

Srsal Srpnit

1 0

3 3

3 3

M-Ocarb HB HB

HB HB

HB OH-O2N

0.144(6) 0.045(2)

Srpf

1

3

3

M-Ocarb M-Oaq

HB

XB

0.035(1)

Srpcl

1

3

3

M-Ocarb M-Oaq

HB

XB

0.035(1)

0.629(8) HB(N)

0.260(3) 0.165(1) 0.160(3)

Baonit

1

2

2

M-Ocarb HB

HB

Bapnh2

2

2f

2f

M-Ocarb HB

M-Ocarb M-Oaq

Basal

2

2

2

M-Ocarb HB

M-Ocarb M-Oaq MLM

Bapams Baben

1 2

3 2

2 2

M-Ocarb M-Oaq M-Ocarb M-Oaq

HB M-Ocarb HB

HB(N) OH 3 3 3 O2N

0.837(22) f

HBpartial

0.659(10) 0.420(9) 0.251(3) 0.061(1)

Bapnit

1

3

3

M-Ocarb HB

HB

Baonh2

2

2

2

M-Ocarb M-Oaq

M-Ocarb

0.054(2)

Bapf

2

3

3

M-Ocarb

M-Ocarb

XB

0.035(1)

Bapcl

1

3

3

M-Ocarb M-Oaq

HB

XB

0.034(1)

0.039(2)

a

Dimensionality of network created by metal to O or N bonds. b Dimensionality of network created by all polar noncovalent intermolecular interactions. c See text for description of modifications made. d Main bonding types seen to propagate network. In most cases the three “dimensions” correspond to the three crystallographic axes—but in some cases diagonals to these axes are used. Key: HB = hydrogen bond of OH-O type; HB(N) = hydrogen bond involving NH2; M-Ocarb = metal to carboxylate bridges; M-Oaq = metal to water bridges; MLM = coordination network propagated by metal ions being linked by bonding to more that one substituent of the benzoic acid; XB = halide to haloarene interactions (X = F or Cl). e The standard deviation in the last decimal place is given in parentheses. f Hydrogen bonds do connect parts of the structure in a third dimension, but these go through the channel hydrate zone. The water sites here are only partially occupied, and thus, this interaction has not been counted above, as the supramolecular network is not complete.

density, enthalpy of hydration, M-O bond strength) are not over-riding determinants of solubility. This highlights the essential problem facing those conducting salt selection studies—even with simple and closely related organic species, there seems to be no way of picking the “best” cation. Note, however, that the rankings are not random. There are 24 possible ways to rank the four cations. The entries for four of the nine anions (sal, onit, pcl, and pf—a group that spans from the most to least soluble anions and covers a wide range of organic substituents) are similar, with Mg giving the most soluble salts, Ca the least soluble, and Ba salts more than or equally as soluble as Sr salts. Overall, Mg is most likely to provide the most soluble salt (5 from 9 cases), and Ca is the most likely to provide the least soluble salt (also 5 from 9). However, all the cations provide at least one example of being the least soluble, and all, except Ba, are in some series the most soluble salt. If it is assumed that Mg salts have a bias to high solubility, possibly in line with the Mg ion’s small size and high charge density,5 then why do four series not show this? All three salt quartets where solvent-separated structures were found for the Mg salts are among these four “anomalous” series. A tentative rule for such group 2 benzoic acid salts could thus be as follows. “Mg forms the most soluble salts except when it adopts a solventseparated ion-pair structure.” This leaves pams as a misfit series. A special pleading here might focus on Capams and Srpams. Both have very different structures from the other Ca and Sr species examined, and it could be this that changes the observed solubility rank ordering for pams. A similar discussion can be initiated around low solubility and Ca salts. All five Ca salts which

are the least soluble members of their series have structures based on 1-D coordination networks (groups iv and v above). This hints at discrete or 2-D Ca structures being more soluble than would otherwise be expected. Great caution is required when extrapolating solubility information from these simple metalbased descriptors. This is highlighted by the three isostructural Ca/Sr pairs of structures. In two cases (the group v channel structures pnh2 and sal), the Ca salts are considerably less soluble than the equivalent Sr salts, which makes sense where both structures make the same intermolecular interactions but the Sr-X interactions can be predicted to be weaker than Ca-X. However, in the final isostructural Ca/Sr pair (pams), the opposite is true. Intermolecular Networks. The above traditional metal-centric inorganic chemical treatment of the structures looked at chemical composition and metal coordination networks. This has some success as a structural tool; for example, it allows some grouping and rationalization of a varied set of structures, but it gives less success with respect to predicting or rationalizing solubility. Differences in solubility are often attributed to differences in hydrogen bonding,7,30,31 and so a more holistic approach, examining metal ion interactions, hydrogen bonding, and halide interactions in combination, was thus attempted. Tables 5 and 6 give a breakdown of the major intermolecular bonding features observed. It can be seen that solubility does not correspond to the dimensionality of the metal-coordination bond network alone. For the Mg salts, the holistic approach works well. It can be seen from column two of the table that the 1325

dx.doi.org/10.1021/cg101547r |Cryst. Growth Des. 2011, 11, 1318–1327

Crystal Growth & Design three most soluble Mg salts have 2-D intermolecular networks and the remaining, less soluble, salts all have 3-D networks. For the other metals, the amine containing derivatives disrupt the pattern. For Ca, Sr, and Ba, ignoring the amine derivatives also gives a consistent pattern of 2-D intermolecular networks being more soluble than 3-D networks. In order to explain all the rankings, including those of the amine derivatives, a lot of manipulation is needed. This can be attempted, but we freely acknowledge that it is somewhat tenuous. For the Sr salts, the pattern of 2-D networks being more soluble than 3-D networks is broken by two salts, Srpnh2 and Srpams, both of which have a para-amino substituent. In both cases, amine based hydrogen bonding helps propagate the network. Such interactions are known to be weaker than OH based hydrogen bonding.31 If amine based hydrogen bonding is ignored, then the most soluble Sr salt is a 1-D network, the four next most soluble are 2-D networks, and the four least soluble are 3-D networks. For both the Ca and Ba salts, similar observations can be made regarding the amino substituted salts. Applying the same argument about amine hydrogen bonding as used above gives Ca and Ba salt ordering, with 2-D structures generally more soluble than 3-D structures. However, there are still misfits, namely the ortho-amino species Caonh2 and Baonh2—both of which are less soluble than their 2-D networks might indicate. There are several obvious arguments against disregarding the amine hydrogen bonds. First, why discard these but retain other weak interactions? In this data set, these include OH 3 3 3 O2N and halide to arene interactions.32 The only counterargument is empirical—this approach gives the best fit to observed solubility. A second point is that while disregarding amine hydrogen bonding gives a better fit between network dimensionality and order of solubility for the Sr, Ba, and Ca salts, it does in fact give a poorer fit for the Mg salts.

’ CONCLUSIONS This work provides a structural database of salts relevant to the pharmaceutical industry—with each salt chosen so as to be systematically chemically and structurally related to the next. The structures of the Mg, Ca, and Sr salts were found to group into a small number of classes, while the Ba salt structures were more variable. Phase specific aqueous solubility data is also presented and tied to the structural data. Taken together, these measurements provide a unique resource for studying structure-property relationships in organic salts.33 For others seeking to emulate the approach of building systematic structural databases, we would highlight the difficulty in obtaining good quality single crystals for any complete group of structures, even where the compounds of interest are chemically simple, and also highlight the problems of correctly measuring phase specific physicochemical data. We present this data with the expectation that it will be used by ourselves and others as the challenging basis for further studies (i.e., lattice energy calculations, thermal studies, chemometrical analysis) on structure-property relationships. We suggest that finding a predictor or predictors that successfully reproduce the rank order solubilities of the 36 salts discussed would be a major advance in our understanding of how array structure influences macroscopic physicochemical properties. Whilst we cannot claim to hold the key to this problem, our initial analysis of the database does raise some salient points. Notably, and contrary to the extant literature, there appears to be little or no correlation of

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solubility with the polarity of the organic ions, with the identity or number of the organic substituents, with cation size, or with hydration state. The intrinsic solubility of the free acid does appear to have a major contribution to make to the salt solubility and so, to some extent, does the identity of the metal cation. Thus, it was observed that Mg tended to form the most soluble salts while Ca tended to give the lowest solubility salts. However, consideration of these two “chemical identity” parameters alone does not successfully predict salt solubility. Adding information about the solid-state structures adopted improves the situation somewhat. Thus, unlike previous studies,12a we do observe a marked qualitative dependence of solubility on solid-state structure. At a crude level, it was found that “Mg forms the most soluble salts, except when it adopts a solvent-separated ion-pair structure”. This is a statement that fits eight of the nine carboxylate anion based series. On a more detailed level, it was shown that considering other noncovalent intermolecular contacts further refines the effectiveness of the model. Considering the dimensionality of the whole intermolecular network gave a good match to solubility rank order for the Mg salts. A similar approach was less successful for the Ca, Sr, and Ba salts, where the amine containing anions disrupted the observed relationship, perhaps indicating a significant difference between N-H centered hydrogen bonding and O-H centered bonding.

’ ASSOCIATED CONTENT

bS

Supporting Information. Details of single crystal characterizations, including cif files and details of powder diffraction measurements on solids both before and after slurry experiments and microanalysis data. This information is available free of charge via the Internet at http://pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*Fax: (þ44) 141-548-4822. E-mail: [email protected].

’ ACKNOWLEDGMENT We gratefully acknowledge WestCHEM for funding a studentship (J.-B.A.). Thanks are also due to the EPSRC National Crystallography Service at the University of Southampton for data collection on Baben, Mgpf, Bapf, and Capcl and to the CCLRC for a beamtime award at the Daresbury Synchrotron Radiation Source. A.R.K. also thanks both Dr. Norman Shankland (Crystalografx Ltd.) and Prof. Alan Cooper (University of Glasgow) for patiently dealing with many questions regarding the measurement and meaning of solubility data. ’ REFERENCES (1) (a) Stahl, P. H., Wermuth, C. G., Eds. Handbook of Pharmaceutical Salts. Properties, Selection and Uses; Wiley-VCH: Zurich, 2002. (b) Gould, P. L. Int. J. Pharm. 1986, 33, 201. (c) Serajuddin, A. T. M. Adv. Drug Delivery Rev. 2007, 59, 603. (2) Hursthouse, M. B. Cryst. Rev. 2004, 10, 85. (3) Shefter, E.; Higuchi, T. J. Pharm. Sci. 1963, 52, 781. (4) (a) Forbes, R. T.; York, P.; Davidson, J. R. Int. J. Pharm. 1995, 126, 199. (b) Rubino, J. T. J. Pharm. Sci. 1989, 78, 485. (5) (a) Anderson, B. D.; Flora, K. P. In The Practice of Medicinal Chemistry; Wermuth, C. G., Ed.; Academic Press: London, 1996; pp 739-754. (b) Chowhan, Z. T. J. Pharm. Sci. 1978, 67, 1257. (c) Anderson, B. D.; Conradi, R. A. J. Pharm. Sci. 1985, 74, 815. 1326

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Crystal Growth & Design (6) Agharker, S.; Lindenbaum, S.; Higuchi, T. J. Pharm. Sci. 1976, 65, 747. (7) Parshad, H.; Frydenvang, K.; Liljefors, T.; Sorensen, H. O.; Larsen, C. Int. J. Pharm. 2004, 269, 157. (8) (a) Lewis, G. R.; Steele, G.; McBride, L.; Florence, A. J.; Kennedy, A. R.; Shankland, N.; David, W. I. F.; Shankland, K.; Teat, S. J. Cryst. Growth Des. 2005, 5, 427. (b) Callear, S. K.; Hursthouse, M. B.; Threlfall, T. L. CrystEngComm 2010, 12, 898. (9) Allen, F. H. Acta Cystallogr., B 2002, 58, 380. (10) For examples see:(a) Yadav, M. R.; Shaikh, A. R.; Ganesan, V.; Giridhar, R.; Chadha, R. J. Pharm. Sci. 2008, 97, 2637. (b) Llin
as, A.; Burley, J. C.; Box, K. J.; Glen, R. C.; Goodman, J. M. J. Med. Chem. 2007, 50, 979. (c) Galcera, J.; Molins, E. Cryst. Growth Des. 2009, 9, 327. (11) Nielsen, A. B.; Frydenvang, K.; Liljefors, T.; Buur, A.; Larsen, C. Eur. J. Pharm. Sci. 2005, 24, 85. (12) (a) Black, S. N.; Collier, E. A.; Davey, R. J.; Roberts, R. J. J. Pharm. Sci. 2007, 96, 1053. (b) Collier, E. A.; Davey, R. J.; Black, S. N.; Roberts, R. J. Acta Crystallogr., B 2006, 62, 498. (c) Chency, M. L.; Shan, N.; Healey, E. R.; Hanna, M.; Wojtas, L.; Zaworotko, M. J.; Sava, V.; Song, S. J.; Sanchez-Ramos, J. R. Cryst. Growth Des. 2010, 10, 394. (13) In our hands, this model approach has previously been used to successfully probe structure in sulfonated azo colorants, see:(a) Kennedy, A. R.; Kirkhouse, J. B. A.; McCarney, K. M.; Puissegur, O.; Smith, W. E.; Staunton, E.; Teat, S. J.; Cherryman, J. C.; James, R. Chem.—Eur. J. 2004, 10, 4606. (b) Kennedy, A. R.; Kirkhouse, J. B. A.; Whyte, L. Inorg. Chem. 2006, 45, 2965. (c) Kennedy, A. R.; Andrikopoulos, P. C.; Arlin, J.-B.; Armstrong, D. R.; Duxbury, N.; Graham, D. V.; Kirkhouse, J. B. A. Chem.—Eur. J. 2009, 15, 9494. (14) XEQMOV, XEQMUB & XEQNAI Caonh2, Sronh2 & Baonh2;Murugavel, R.; Karambelkar, V. V.; Anantharaman, G.; Walawalkar, M. G. Inorg. Chem. 2000, 39, 1381. (15) JEDCK1 Caben;Senkovska, I.; Thewalt, U. Acta Crystallogr., C 2005, 61, m448. (16) FUVCII Capf;Karipides, A.; McKinney, C.; Peiffer, K. Acta Crystallogr., C 1988, 44, 46. (17) UNATEI Mgpnh2;Krishnamurthy, D.; Sathiyendiran, M.; Murugavel, R. Proc. Ind. Acad. Sci.: Chem. Sci. 2000, 112, 395. (18) BIFDOJ Srpnh2;Amiraslanov, I. R.; Musaev, V. N.; Mamedov, Kh. S. Zh. Struk. Khim. 1982, 23, 114. (19) VAXJAF Mgpams;Cole, L. B.; Holt, E. M. Inorg. Chim. Acta 1989, 160, 195. (20) VAXHUX01 Mgsal;Drake, S. R.; Sanderson, K. D.; Hursthouse, M. B.; Malik, K. M. A. Inorg. Chem. 1993, 32, 1041. (21) CASALA01 & CSSALB Casal & Srsal;Debuyst, R.; Dejehet, F.; Dekandelaer, M.-C.; Declercq, J. P.; van Meerssche, M. J. Chim. Phys. Phys.-Chim. Biol. 1979, 76, 1117. (22) XITDUZ Mgonh2;Wiesbrock, F.; Schier, A.; Schmidbaur, H. Z. Naturforsch., B: Chem. Sci. 2002, 57, 251. (23) Cernik, R. J.; Clegg, W.; Catlow, C. R. A.; Bushnell Wye, G.; Flaherty, J. V.; Greaves, G. W.; Burrows, I.; Taylor, D. J.; Teat, S. J.; Hamichi, M. J. Synchrotron Radiat. 1997, 4, 279. (24) Sheldrick, G. M. Acta Crystallogr. 2008, A64, 112. (25) The acids which gave incomplete groups of structures were p-bromobenzoic acid, p-hydroxybenzoic acid, p-methylbenzoic acid, ofluorobenzoic acid, o-chlorobenzoic acid, o-methylbenzoic acid, and dihydroxybenzoic acid. The principle reason for failure was that the halo and methyl substituted acid salts tended to form multiple crystals (typically, platelike crystals that stacked together in a misaligned fashion). (26) Manojlovic, L.; Speakman, J. C. J. Chem. Soc. A 1967, 971. (27) See for instance:(a) Julian, M. O.; Day, V. W.; Hoard, J. L. Inorg. Chem. 1973, 12, 1754. (b) Bach, I.; Kumberger, O.; Schmidbaur, H. Chem. Ber. 1990, 123, 2267. (c) Nicolis, I.; Coleman, A. W.; Charpin, P.; de Rango, C. Angew. Chem. 1995, 34, 2381. (28) C^ote, A. P.; Shimizu, G. K. H. Chem.—Eur. J. 2003, 9, 5361. (29) Galcera, J.; Molins, E. Cryst. Growth Des. 2009, 9, 327. (30) Ghasemi, J.; Saaidpour, S. Chem. Pharm. Bull. 2007, 55, 669.

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(31) The classic case of different solubilities in polymorphs of Ritonavir is often described with respect to differences in O-H 3 3 3 O versus N-H 3 3 3 O bonding, see:(a) Chemburkar, S. R.; Bauer, J.; Deming, K.; Spiwek, H.; Patel, K.; Morris, J.; Henry, R.; Spanton, S.; Dziki, W.; Porter, W.; Quick, J.; Bauer, P.; Donaubauer, J.; Narayanan, B. A.; Soldani, M.; Riley, D.; MaFarland, K. Org. Process Res. Dev. 2000, 4, 413. (b) Bauer, J.; Spanton, S.; Henry, R.; Quick, J.; Dziki, W.; Porter, W.; Morris, J. Pharm. Res. 2001, 18, 859. (32) (a) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Acc. Chem. Res. 2005, 38, 386. (b) Price, S. L.; Stone, A. J.; Rowland, L. R. S.; Thornley, A. E. J. Am. Chem. Soc. 1994, 116, 4910. (33) Some of the problems with the availability of suitable solubility data sets are summarized in:Llin
as, A.; Glen, R. C.; Goodman, J. M. J. Chem. Inf. Model. 2008, 48, 1289. (34) Yalkowsky, S. H., He, Y., Eds. Handbook of Aqueous Solubility Data; CRC Press: Boca Raton, FL, 2003.

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