An Examination of Water Vapor Sorption by Multicomponent Crystalline and Amorphous Solids and Its Effects on Their Solid-State Properties

Journal of Pharmaceutical Sciences 108 (2019) 1061-1080

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General Commentary

An Examination of Water Vapor Sorption by Multicomponent Crystalline and Amorphous Solids and Its Effects on Their Solid-State Properties Ann Newman 1, *, George Zografi 2 1 2

Seventh Street Development Group, Kure Beach, North Carolina 28449 School of Pharmacy, University of Wisconsin-Madison, Madison, Wisconsin 53706

a r t i c l e i n f o

a b s t r a c t

Article history: Received 20 September 2018 Revised 23 October 2018 Accepted 24 October 2018 Available online 1 November 2018

This commentary critically evaluated the unique effects of water vapor sorption by multicomponent solid forms of active pharmaceutical ingredients (APIs), and its effects on their physical and chemical properties. Such multicomponent forms include the following: (1) crystalline salts and cocrystals, and (2) amorphous salts, coamorphous mixtures, and amorphous solid dispersions (ASDs). These solid forms are commonly used to increase the solubility, dissolution, and bioavailability of poorly soluble APIs. To achieve this increase, selected counterions or coformers exhibit much greater polarity, and have a tendency to enhance water vapor sorption, leading to possible instabilities. Such instabilities include salt disproportionation, cocrystal dissociation, and phase separation and crystallization from amorphous forms. Regarding crystalline multicomponent systems, significant instabilities arise on account of deliquescence or crystal hydrate formation. Such behavior often follows water-induced salt disproportionation or cocrystal dissociation. Regarding amorphous salts, coamorphous mixtures, and ASDs, we see the importance of absorbed water as a disrupter of API-coformer interactions and as a plasticizer in bringing about subsequent phase separation and crystallization. In preparing multicomponent solid forms, it is important to measure the water vapor sorption isotherm of the counterion or coformer to better understand the mode by which water is sorbed, and to anticipate and correct possible instabilities. © 2019 American Pharmacists Association®. Published by Elsevier Inc. All rights reserved.

Keywords: amorphous solid dispersion(s) cocrystal(s) deliquescence hydrate(s) hydration phase transformation(s) physical stability salts salt disproportionation water sorption

Introduction It is well established that the physical and chemical stability of solids, in the form of active pharmaceutical ingredients (APIs), excipients, and formulations, very often decreases upon exposure to environments containing water vapor. For this reason, measurement of the “equilibrium” amount of water vapor sorbed by solids as a function of relative humidity (RH) and at constant temperature is a routine part of the drug development process.1,2 In addition, because water vapor molecules can interact by different modalities, it is important to determine the dominant molecular mechanisms by which such sorbed water might interact with the solid, and possibly affect its structure and dynamics in such a way as to cause physical and chemical change. In this article, we wish to focus attention on water vapor sorption by binary “multicomponent” solids, where the components form a miscible mixture at the molecular level through various types of

* Correspondence to: Ann Newman (Telephone: þ1-765-650-4462). E-mail address: [email protected] (A. Newman).

intermolecular interaction mechanisms. In this group of solids, we include the following: (1) crystalline salts (API and counterions) and cocrystals (API and small molecule coformers), and (2) amorphous systems (amorphous API salts, coamorphous API [small molecule coformers], and amorphous solid dispersions [ASDs] containing API and polymers). We primarily are interested in knowing how such interactions between a particular API and different counterions or coformers might affect water vapor sorption, and the tendencies for such water vapor sorption to alter intermolecular interactions between components to bring about physical and chemical transformations. Physical transformations would include deliquescence, crystal hydrate formation, disproportionation of salts, dissociation of cocrystals, and phase separation of coamorphous and amorphous dispersion systems to form separated amorphous phases of each component, and possible API crystallization. We will begin by briefly reviewing various modes of water vapor sorption by solids, and then focus on examples of such behavior with multicomponent systems. We also will focus attention on more complex formulation systems, containing miscible multicomponent solids along with other formulation ingredients, where water vapor sorption gives rise to various instabilities involving API and excipients. The goal of this

https://doi.org/10.1016/j.xphs.2018.10.038 0022-3549/© 2019 American Pharmacists Association®. Published by Elsevier Inc. All rights reserved.

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article is to develop a better conceptual understanding of the relationship between water vapor sorption and the physical and chemical properties of various multicomponent solids, using selected pertinent literature for our analysis. Water Vapor Sorption by Solids Overview of Adsorption and Absorption The term “sorption” generally refers to interactions of water vapor with solids by 2 possible broad mechanisms: (1) adsorption where the interactions occur only between water and molecules located at the surface and (2) absorption where water molecules are also able to penetrate into the bulk phase of the solid. Scheme 1 illustrates both of these broad categories of sorption along with a further breakdown of the various more specific mechanisms possible under each category. Adsorption leads to an alteration in the thermodynamic and kinetic properties of molecules at the surface, while absorption also can have significant effects on such properties in the bulk phase. The overall tendency for both adsorption and absorption of water molecules by a solid, under a given set of temperature and RH conditions, can be described thermodynamically at any temperature, T, by the free energy change, DGs, which in turn can be expressed by the enthalpy change, DHs, and entropy change, DSs, from the Gibbs-Helmholtz equation:

DGs ¼ DHs  TDSs

(1)

Because the process of sorption involves transferring a water molecule from the free vapor state to one where interaction with the solid has occurred, and hence introducing decreased degrees of freedom for water molecules, we ordinarily would expect DSs to be extensively negative. Hence, to produce an overall negative free energy change under the ambient temperatures normally encountered, there must be enough intermolecular interaction to produce a negative DHs or a more positive DSs associated with changes in the surface and bulk solid-state structure caused by interactions with water, for example, more disorder. On this thermodynamic basis, the sorption of water vapor by a solid at ambient temperatures through noncovalent interaction should be greater, the greater the ability of the solid to hydrogen bond with water. Water Vapor Sorption by Crystalline Solids Physisorption As introduced above in Scheme 1, a crystalline solid generally can only adsorb water on its surface unless water molecules are

able to penetrate the crystalline lattice to form stoichiometric or nonstoichiometric hydrates, or to deliquesce. In such cases, often referred to as physisorption, it appears that the amount of water adsorbed at RHs below any deliquescence is limited to the equivalent of just a few “molecular layers.”3 The word, equivalent, is emphasized, because unless the surface is energetically homogeneous, the layers of water molecules most likely will be distributed at different energetic sites, such as defects. From estimates of the cross-sectional area of a water molecule, and assuming closest packing of molecules to form each monolayer on the available surface, we can estimate that the equivalent of, for example, 3 molecular layers per square meter (m2) of surface would amount to no more than about 0.1% w/w water uptake. We would expect the first layer of water molecules, which directly interact with the solid to have a molecular orientation and general structure determined in large part by the mechanism and strength of the solid-water interaction. Such a surface water structure apparently then can attract other water molecules to form additional “layers,” where these water molecules do not yet assume the exact structure of bulk liquid water capable of dissolving the solid. Deliquescence As shown in Scheme 1, solids with significant polarity, for example, salts and sugars, exposed to higher RH, above that of a saturated aqueous solution (RHo) of the water-soluble solid, deliquescence with the formation of liquid water on the surface capable of dissolving the solid.4 Water vapor sorption, liquefaction, and dissolution of the solid at RH values at and greater than RHo will continue to increase with time until the solid phase creating the saturated solution is no longer present. Such behavior for the crystalline salt, ranitidine HCl, for example, occurs just above 75% RH, as illustrated in Figure 1.5 Note the reversibility of the equilibrium point RHo after successive cycles of sorption/desorption, while the values of water uptake above RHo are slightly different on account of the time-dependent, nonequilibrium uptake of water above RHo. Capillary Condensation For highly porous crystalline solids, with pore dimensions less than 100 nm, where the pores have a very high degree of curvature, an extra-thermodynamic effect characterized by the Kelvin equation can lead to premature (less than 100% RH) liquefaction of water in the form of “capillary condensation” and possible solid dissolution.6 Generally, this is not a major factor in the effects of water vapor sorption on the properties of pharmaceutical solids, but it should be kept in mind as a possible factor when dealing with highly porous and granular water-soluble materials at relatively

Modes of Water Vapor SorpƟon

AdsorpƟon

AbsorpƟon

(surface)

(bulk)

LiquefacƟon

Deliquescence (water soluble crystals)

Crystal Hydrate

DissoluƟon in Amorphous Solids

Capillary CondensaƟon (porous solids)

Scheme 1. Possible modes of water vapor sorption by crystalline and amorphous solids.

Figure 1. Water vapor sorption of ranitidine HCl indicating the critical deliquescence point (RHo). Reproduced with permission from Salameh and Taylor.5

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high RH, and where significant amounts of such confined space exists. Crystal Hydrates The formation of crystal hydrates by exposure to water vapor generally occurs with molecules that undergo solution-mediated crystallization to form hydrates in the presence of various amounts of liquid water at a particular water activity, aw. From a thermodynamic perspective, the same crystal hydrate should form when the solid is at an RH of comparable water activity (RH, expressed as a fraction, is approximately equal to water activity). However, this generally occurs at lower rates of transformation and at a different critical water activity controlled by the rates of solidstate diffusion and crystal nucleation in the bulk phase. An interesting example of such behavior is the water vapor sorption by anhydrous sodium naproxen as a function of RH and temperature.7 Here, as described in the phase diagram depicted in Figure 2, one can obtain a monohydrate and tetrahydrate, as well as 2 polymorphic forms of the dihydrate, depending on the RH and temperature. However, water vapor sorption studies at 25 C (Fig. 3), starting with anhydrous crystalline sodium naproxen, indicate some complexities, presumably related to the kinetics of water vapor sorption. Here, for example, we see that the first exposure of water vapor to the anhydrous salt (first sorption) shows no evidence of the monohydrate, as obtained by solution-mediated crystallization, but rather goes directly to the dihydrate (one of the polymorphic dihydrates) mentioned above. Increasing RH indicates a sorption profile consistent with the formation of the tetrahydrate. Note that the first desorption profile (first desorption) indicates the appearance of the tetrahydrate, dihydrate, and monohydrate forms with significant hysteresis upon desorption, most likely, again, due to the kinetics of the process. Presumably, the nucleation of the monohydrate at lower RH occurred only after the formation of the other hydrates. Interestingly, a second sequence of exposure, starting with the monohydrate (second sorption), indicates significant differences in the sorption profiles from the first sorption sequence, but very similar desorption profiles (second desorption). Understanding the thermodynamic basis for the formation of crystal hydrates of varying stoichiometry and physical and chemical properties is a critical concern when various crystalline solids

Figure 2. A three-dimensional phase diagram indicating various hydrates of sodium naproxen based on water sorption data: monohydrate (MH), dihydrate (DH), and tetrahydrate (TH). Reproduced with permission from Raijada et al.7

Figure 3. Water vapor sorption/desorption of sodium naproxen at 25 C, starting with the anhydrous form (AH). Theoretical water content % w/w: monohydrate (MH) 6.7%, dihydrate (DH) 12.5%, and tetrahydrate (TH) 22.2%. Reproduced with permission from Raijada et al.7

are exposed to water at a given water activity (RH). Key to such behavior is the propensity of water to hydrogen bond with various functional groups of the solid at isolated sites within the crystal lattice, or to form hydrogen bonded water clusters within channels formed by the geometric properties of the organic molecules.8 With regard to isolated sites, in addition to hydrogen bonding between electron donor and acceptor groups, the strong hydration spheres of metallic counterions associated with the lattice, for example, Naþ, Caþþ, also can play an important role. The relationship between the sizes of any lattice channels relative to the type of water clusters that might occur is critical in determining the stability of such hydrates upon changes in RH or in temperature. Such channels are generally, but not always, associated with the formation of nonstoichiometric hydrates and tendencies to form dehydrated hydrates or amorphous forms upon dehydration. Water Vapor Sorption by Amorphous Solids As indicated in Scheme 1, water vapor molecules generally penetrate the surface of amorphous solids and interact with molecules in the bulk phase primarily via hydrogen bonding and on account of the greater free volume (lower density) of the amorphous form relative to that of the corresponding crystal.1,2 To illustrate the impact of the free volume on tendencies for sorption, compare in Figure 4 the sorption of water by the free acid form of indomethacin in the highly ordered crystalline form (density of 1.38 g/cm3) versus that of a less dense amorphous glass, formed by quench cooling of liquid indomethacin (density of 1.31 g/cm3).9 Note also the much less sorption of an amorphous form of indomethacin prepared by vapor deposition at low temperatures. In this case, the process of forming the glass from the vapor state produces a reduction in density of about 1.4% and hence greater resistance to the penetration of water molecules into the bulk phase. Generally, water dissolving into the amorphous solid disrupts the intermolecular forces in the solid and increases the overall free volume of the system, thereby increasing diffusional molecular mobility through its so-called “plasticizing” effects.2 This is reflected as illustrated in Figure 5, as a reduced glass transition temperature (Tg) as a function of water content, given that the value of Tg reflects the temperature at which the diffusional relaxation time and viscosity for the amorphous solid undergo significant change.10 Although what just has been discussed is generally true, that is, sorbed water acts as a plasticizer and reduces the glass transition temperature, it has been shown that at lower RH water vapor is taken up in the

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to selectively enhanced adsorption of water vapor molecules. Such selective adsorption at defect sites, with resulting possible enhanced molecular mobility, often leads to the initiation of physical and chemical changes at the surface.16 Furthermore, the processing of crystalline solids by, for example, milling or drying, can lead to more defects and eventually to partially or fully amorphous solids.17 In such cases, water vapor appears to distribute between crystalline and amorphous regions unequally, favoring the amorphous state. For example, the equilibrium water vapor sorption by crystalline NaCl, ground and unground, on a per unit area basis, is 3 times greater for the ground sample than that for the unground crystal up to about 20% RH, above which crystallization of the disordered regions and a loss of specific surface area occur.18 Water Vapor Sorption by Multicomponent Systems Overview Figure 4. Comparison of water vapor sorption by indomethacin in the form of an ordinary glass, a stable glass prepared by vapor depositions, and the crystal. Reproduced with permission from Dawson et al.9

glassy state (below Tg) of a solid by 2 parallel mechanisms leading to concurrent plasticization and antiplasticization.11,12 Apparently, water vapor sorption occurs in 2 structural regions of a glassy solid, the major domain structure where plasticization occurs, and a higher energy region at the interface of the domains, known as the microstructure, where water acts to reduce free volume and hence to antiplasticize that region. Such antiplasticizing behavior appears related to an increase in relaxation time for secondary JohariGoldstein motions primarily associated with the microstructure.13 It appears that the sorption isotherm of amorphous solids is predicted quantitatively by combining factors influenced by both dissolution into the major domains with increased free volume, and free volume reduction by water entering the microstructure regions.14,15

Influence of Crystal Defects and Partially Amorphous Regions It is well recognized that crystalline solids with well-defined crystal faces will contain a variety of defects where molecules would be expected to be in a higher state of surface energy leading

Figure 5. Plasticizing effects of water vapor sorption on the glass transition temperature (Tg) of poly(vinylpyrrolidone). Reproduced with permission from Oksanen and Zografi.10

We begin by assuming that water vapor sorption by a physical mixture (phase separated) of solid components, under a particular set of RH and temperature conditions, generally is the weighted average of the amount of water vapor sorbed by the individual components under identical conditions.19,20 For example, the total weight of water sorbed per dry weight of solid by a binary physical mixture, Wt, is equal to the sum of the weight of water sorbed per dry weight of the individual components, Wa and Wb, weighted by the weight fraction of each component, fa and fb, so that

Wt ¼ ðWa f a Þ þ ðWb f b Þ

(2)

where fa is equal to 1  fb. That such behavior occurs with physically mixed systems is illustrated in Figure 6 for a number of excipient physical mixtures.20 We further assume that deviations from such a simple mixing rule will occur if the components form a molecularly miscible mixture, with such deviations depending on the type, extent, and strength of interactions occurring between the components. Water Vapor Sorption of Inorganic Crystalline Salts Because the maximum amount of water vapor adsorbed on a typical crystalline surface with moderate specific surface area is extremely small, ~ 0.1% per m2 of surface, such solids generally are considered to be nonhygroscopic, and very few quantitative adsorption studies with crystalline pharmaceutical systems have been reported in the literature. What we do know about water vapor adsorption at a mechanistic level primarily comes from many studies with highly purified and well-characterized crystalline inorganic salts, such as members of the alkali halide family.21-23 As shown in Figure 7, for example, the 001 crystal face of a highly purified crystalline samples of NaCl appears to take up a maximum of about 3-4 molecular layers up to 75% RH where deliquescence occurs.22 Note the distinct break in the plots at the equivalent of 1, 2, and 3 molecular layers, indicating that the surface energetics of the NaCl crystal are relatively homogeneous. As also seen in Figure 7, the amount of water sequentially adsorbed and desorbed at a particular RH is essentially constant, indicating that such adsorption does not appear to have a significant irreversible effect on the bulk properties of the crystal. However, independent experimental measurements, for example, infrared, have shown that, in such a system, adsorbed water at the level of just a few molecular layers accumulate as a “liquid-like layer” which is capable of hydrating Naþ and Cl ions. Such hydration as it continues then increases local ionic motions without bulk dissolution as with deliquescence. In Figure 8, we can see a schematic

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Figure 6. Predictions from Equation 1 (in text) versus experimental water vapor sorption for various mixtures of tablet excipients. Left side: comparison of tablets, granules, and dry blends for one formulation. Right side: comparison of dry blends of different mixtures. Reproduced with permission from Dalton and Hancock.20

representation of the possible state of adsorbed water molecules (white circles) and ions (dark circles) at surface coverages up to 3 molecular layers (below deliquescence), and at ambient temperatures versus cryogenic temperatures.23 Note that under ambient temperatures some ionic movement into the adsorbed layer occurs even in a single molecular layer, with significantly greater ionic accumulation and motions with increasing coverage. On the other hand, no such accumulation appears to occur at any coverage at the lower cryogenic temperature, consistent with less ionic mobility at the lower temperature. Further studies under ambient temperature with other alkali halides indicate differences related to the

preferential solvation of cations.24 In another study, it was shown that LiF, a relatively water-insoluble salt on account of the very small size of the Liþ and F ions, and hence very strong electrostatic attraction, does not show a tendency for increased ionic mobility in the adsorbed layers under ambient temperatures.25 These results suggest that water vapor adsorption on any crystalline surface

Figure 7. Water vapor adsorption (dark circles)/desorption (open circles) on crystalline sodium chloride (001) at 24 C. Reproduced with permission from Foster and Ewing.22

Figure 8. Schematic representation of proposed models for the first few “layers” of adsorbed water vapor on the surface of crystalline NaCl (100) at cryogenic and ambient temperatures. Reproduced with permission from Peters and Ewing.23

A. Newman, G. Zografi / Journal of Pharmaceutical Sciences 108 (2019) 1061-1080

below deliquescence can form a “liquid-like layer” where, if sufficiently hydrated, the ions making up the salt are in a higher state of energy and mobility. Further evidence points to such water vapor adsorption preferentially occurring at high-energy crystal defect sites so that the distribution of water molecules on the surface at such low levels of adsorption can be quite heterogeneous. The possible effects on physical and chemical stability of organic salts at such low levels of water vapor adsorption, particularly for relatively polar organic API and counterion combinations and cocrystals, are worthy of further consideration and analysis (to be discussed below).

100

RH of Saturated SoluƟon (%)

1066

oxalate

95

(+) tartrate (-) tartrate

90

hydrochloride

85

80 methanesulfonate

75

Water Vapor Sorption of Organic Crystalline Salts

0.0

Role of Counterions The major questions of interest with regard to water vapor sorption by organic salts of an API is the role played by a particular counterion relative to others in affecting sorption as well as in affecting various physical and chemical properties. Regardless of the mechanisms involved, we would expect the counterion size and polarity to have some influence on the strength of interaction between API and counterion and some effect on hydration tendencies. For example, as reflected in Table 1, for 3 deliquescent choline halide salts, a significant reduction in RHo occurs in the following order: chloride < bromide < iodide, an order corresponding to that of ionic size of the counterion and the aqueous solubility of the salt.26 In a related study, as shown in Figure 9, the deliquescence of a series of an experimental API salt with different counterions was observed, likewise, to be inversely proportional to the aqueous solubility of the salts.27 Hence, during a salt selection process for this API in early development, the advantages of using the most water-soluble salt, the methanesulfonate, to enhance oral bioavailability were offset by a significant decrease in RHo and, hence, an increased tendency for deliquescence. In another study of counterion effects on water vapor sorption characteristics, carried out during a process of salt selection for an experimental API, BMS-180431, significant differences between the various counterions were demonstrated by Morris et al.28 Here, 7 salts of the API (sodium, potassium, calcium, zinc, magnesium, arginine, and lysine) were formed as crystals and subjected to storage at a number of RHs, as shown in Table 2. Also included in Table 2 are initial water contents of the salts after exposure to ambient humidity and temperature and unspecified drying conditions. In all cases, the amount of water present initially and sorbed at various RH values appears to be greater than we would expect for only adsorption to a crystalline solid surface. In addition, there was no evidence that deliquescence had occurred, or that water sorption was due to the presence of amorphous regions. Consequently, there appears to be a strong possibility that all the salts may have formed various types of crystal hydrates upon exposure to different RH values. In Table 3, the calculated expected weight percent of water sorbed at various water-to-salt molar ratios, from a monohydrate to a hexahydrate, appears to be in the range experimentally observed at various RH in this study. Interestingly, the loss of water upon heating in the differential scanning

Table 1 Aqueous Solubility and Critical Deliquescence Point (RHo) of Various Choline Halides26 Salt

Apparent Solubility (mol/L)

RHo (%)

Choline chloride Choline bromide Choline iodide

32.2 25.9 14.0

23 41 73

0.2

0.4

0.6

0.8

1.0

1.2

Solubility (M) Figure 9. Relationship between the critical deliquescence point (RHo) and the aqueous solubility of various API salts. Adapted from Nicklasson and Nyqvist.27

calorimetry for the 5 metal ion salts occurred in 2 steps, suggesting the possible ability to form multiple crystal hydrates depending on the RH. On the other hand, the more hydrophobic amino acid salts showed less sorption, constant at all RH values, and a single step for thermal dehydration. Further X-ray studies with the magnesium salt indicated a crystal structure change upon exposure to elevated RH, whereas that of the amino acids did not show such a change. The higher values of water sorbed by the amino acid salts, relative to those expected for surface adsorption, and constant at various RH, suggest that perhaps a loosely held monohydrate might be forming, while in contrast to the metal ion salts, no higher hydrates occur at the higher RH conditions. Minimally, this study has shown that the different properties of the various counterions used in this study have significantly different effects on the properties of the API salt former, and that the more hydrating metal ions promote greater water vapor sorption and greater possibility of hydrate formation than the more nonpolar amino acids. Crystal Hydrate Formation of Salts In regard to counterion effects, an interesting hydrate-forming API is the oral hypoglycemic drug, sitagliptin, marketed as the monohydrate of the phosphate salt, which, as shown by Tieger et al.,29 also is produced as the L-tartrate salt, which can be crystallized as a hemihydrate with at least 4 polymorphic forms. Water sorption-desorption analysis, thermal analysis, and variable humidity X-ray powder diffraction and single crystal structure determination indicate that the water involved in forming the hydrate does not trigger any phase transformations. Rather, water molecules are located in lattice channels, consisting of infinite sheets formed by hydrogen tartrate anions linked by parallel chains of hydrogen bonded water. Removal of the water by heat, or lowering of the RH to dryness, produces isostructural dehydrates. The importance of hydration by metallic counterions in the Table 2 Moisture Content for Storage at Various Relative Humidities for BMS-108043128 Salt

% Moisture by Dry Weight Initial

33% RH

52% RH

75% RH

Sodium Potassium Calcium Zinc Magnesium Arginine Lysine

4.4 9.5 5.0 4.4 13.8 2.8 3.2

4.4 10.6 10.4 2.9 13.5 3.0 3.6

16.0 17.1 13.4 6.5 14.0 3.2 4.0

ND ND 16.1 8.3 14.9 3.8 6.0

ND, not determined.

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Table 3 Calculated Percent Water Sorbed for Various Stoichiometric Hydrates for BMS-180431 Salts Salt

Monohydrate

Dihydrate

Trihydrate

Tetrahydrate

Pentahydrate

Hexahydrate

Free base Sodium Potassium Calcium Zinc Magnesium Arginine Lysine

3.8 3.6 3.5 3.5 3.3 3.6 2.8 2.9

7.3 7.0 6.8 6.8 6.5 7.0 5.4 5.6

10.6 10.1 9.8 9.8 9.4 10.1 7.9 8.2

13.6 13.1 12.7 12.7 12.1 13 10.2 10.7

16.5 15.8 15.4 15.3 14.7 15.8 12.5 13.0

19.1 18.4 17.9 17.9 17.1 18.3 15.6 15.2

formation of crystal hydrates also was demonstrated in a comprehensive series of studies with the drug, nedocromil and its salts. Here, for example, nedocromil sodium exists as a monohydrate, trihydrate, and heptahemihydrate, with distinctly different crystal structures depending on the RH.30 In this system, the role of counterion hydration is very significant with sodium ions and water closely associated in crystal lattice channels. It is also interesting to note that nedocromil, which is a dicarboxylic acid, has the ability to form a wide range of other structurally distinct hydrates with a variety of bivalent counterions, including calcium, manganese, cobalt, magnesium, and zinc, where hydration of these cations plays an important role.31-33 Counterion Effects on Sorption Isotherms As another example where counterion properties affect water vapor sorption of crystalline solids, we can examine the comprehensive work of Guerrieri et al.,34 who prepared a number of salts of the basic drug, procaine, and measured water vapor sorption at 25 C and 50 C. Because one of the important properties of the salts evaluated in this study was the effect of sorbed water on chemical stability at 50 C (to be discussed below), Figure 10 provides a comparison of water vapor sorption for 10 of the procaine salts at 50 C. Here, the amount of sorption is presented as the weight change of water per m2 of solid surface area versus RH to normalize all data to a per unit area basis. Extensive studies showed that all salts were initially in a crystalline state and that significant absorption due to the presence of amorphous regions was not a likely factor. Note the wide variation in the shape of the isotherms among the various salts and the marked differences for water sorbed at any particular RH. Typically, for a single deliquescent component that equilibrates rapidly at RHo, the amount of water sorbed shows very little adsorption below RHo followed by an abrupt increase at and above RHo, as shown for ranitidine HCl in Figure 1. Note in Figure 10

Figure 10. Water vapor sorption isotherms at 50 C for crystalline salts of procaine. Reproduced with permission from Guerrieri and Taylor.34

that 3 salts, the estylate, hydrochloride, and oxalate, appear to exhibit this characteristic shape, while all other salts appear to show a more gradual increase in sorption with increasing RH. Given that the authors indicated that at least 6 salts were deliquescent (RHo values were not reported for all salts), one would assume that for some salts the deliquescence rates were slower than normally expected. A possible factor might be a slower rate of dissolution upon deliquescence,4 or, perhaps, other complicating factors such as concurrent chemical degradation.

Effects of pH The solubility of saturated solutions of each salt at 50 C is shown in Table 4. In contrast to distinct correlations between aqueous solubility and the values of RHo, shown above for other systems, here we see behavior that is more complex. For example, although the hydrochloride salt is the most water-soluble salt at 50 C, it appears to take up less water, with a high RHo, compared to the other soluble salts, with the exception of the poorly soluble oxalate salt. However, if we compare the various alkyl and aryl sulfonic acid salts, we do see a rough correlation between water vapor sorption and solubility, with decreasing sorption at a particular RH as we go from the more soluble mesylate salt to the less soluble napsylate and tosylate salts. Other comparisons are difficult to make on account of the complex shapes of the various isotherms, with significant and continuous sorption at all RH values and no clean-cut RHo. One other source of complexity in comparing the water vapor sorption of these salts is the possibility that the different counterions produce different degrees of ionization of procaine salts in the sorbed layers and, therefore, different tendencies for the amount of water sorbed and the effects of different salts on procaine properties. One indication of such differing acid-base behavior is the pH of the various saturated solutions, shown in Table 4, where we can see that pH values vary widely. This is an important parameter to consider because the pH of a saturated solution appears to correlate somewhat with what might be expected to be a microenvironmental pH of the sorbed layers of water at the crystalline surface.35 To assess the possible role of salt solubility, levels of water sorption, and the pH produced in the saturated sorbed layers of water in affecting solid properties, Guerrieri et al.34 carried out a very comprehensive analysis of solid-state procaine hydrolysis at 50 C. The basic hypothesis for this analysis was that the solid-state chemical degradation rates at a given RH for each salt arose from a combination of (1) amount of water sorbed; (2) the aqueous equilibrium solubility; and (3) the pH of the sorbed layer of hydrated salt. Thus, water sorbed by the solid supposedly creates a “saturated solution” of each salt at a given pH, leading to a pseudo zero-order degradation rate, as expected for a saturated solution of a molecule that undergoes first-order solution degradation as in a suspension. Using solution kinetic data as a function of pH and temperature, they were able to estimate a predicted rate constant, Kpred(s) for a given amount of water sorption, taking into account

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Table 4 Aqueous Solubility, pH of a Saturated Solution, Storage RH, and Number of Equivalent Monolayers Sorbed for Various Procaine Salts at 50 C34 Salt

Solubility in Water at 50 C (mol/L)

pH of Saturated Solution at 50 C

Storage RH at 50 C

No. of Sorbed Monolayers

Napsylate Oxalate Tosylate Phosphate Citrate Besylate Estylate Mesylate Bisulfate Formate HCl

0.049 0.238 0.358 0.870 1.20 2.14 3.00 3.29 3.55 3.75 3.84

6.47 3.44 5.53 4.52 3.72 5.98 5.91 4.69 1.26 6.42 4.76

74 74 74 69 69 74 45 31 31 21 74

15 2.1 6 29 10 13 11 32 15 6 0.9

the equilibrium solubility and pH of a saturated solution, shown in Table 4. To test the hypothesis at relatively low amounts of sorption, presumably, below any RHo, samples were stored at 50 C, at a given RH. Table 4 indicates the storage RH and the calculated number of equivalent molecular layers of sorbed water for the various salts. Note that the number of sorbed layers ranged from about 1 monolayer for the HCl salt to 32 for the mesylate salt. Because, except for the oxalate salt with an equivalent number of monolayers equal to 2.1, all other values ranged from 6 to 29, with what would appear to be more than the normally expected amount for simple physisorption, that is, about 3-4 equivalent molecular layers. However, this still represents a very limited amount of sorbed water. From estimates of initial zero-order kinetics for the salts at various storage RHs, taking into account the amount of water sorbed, the aqueous solubility, and the pH of a saturated solution, it was possible to calculate the solid-state rate constant Kobs(s) for the various salts and to show very consistent agreement with the corresponding values predicted from solution data. For example, studies with procaine besylate gave a Kpred(s) of 1.23  106 gsalt 6 gsalt hydrated/total gsalt hydrated/total gsalt h, and a Kobs(s) of 1.33  10 h (comparisons of all other salts can be found in Guerrieri and Taylor36). In view of the relatively small amount of water present, the authors suggested that the type of limited hydration shown in Figures 7 and 8 for NaCl, water could be playing a role as a “liquidlike layer” in these systems. One possible aspect of this, which would amplify the ability of small amounts of sorbed water to act as a “solvent,” would be the known tendency for initially adsorbed water molecules to accumulate selectively at high-energy surface sites, such as defects or small amounts of amorphous material, and to concentrate the number of associated water molecules. The important point of this work, however, is that the sorbed layers of water, even at very small amounts, appear to produce an environment where the solubility and microenvironmental pH of a salt can play a major role in determining an important solid-state property. The fact that solid-state chemical degradation of procaine salts was quantitatively predicted at very low levels of water sorption, suggests the need for further work with, perhaps, very well-defined single crystal organic salts, where the possibility that “liquid-like layers,” as described above for the alkali halides, can be more systematically evaluated. Of particular interest would be understanding the ability of the sorbed layer to solvate the organic ionic species and promote diffusional mobility.

Salt Disproportionation Another solid-state transformation of salts, which reveals the importance of counterions and sorbed water, is the process of salt disproportionation, whereby salts of weak acids and bases in the solid state undergo a reaction to the corresponding less watersoluble free acid or free base of the ionic API and its counter

ion.36-39 Evaluation of solution equilibria associated with the disproportionation of salts of weak acids or bases to produce their corresponding free base or acid as a function of pH, leads to a pHmax, which is the pH at which the solid phases of the salt and free form are both at equilibrium.40,41 For a salt of a weak base, we can write

pHmax ¼ pKa þ logfSFB =SMS g

(3)

where pKa is the dissociation constant for the conjugate acid of the base, SFB the solubility of the free base, and SMS the solubility of the salt. Correspondingly, the pHmax for the salt of a weak acid is given as:

pHmax ¼ pKa þ logfSMS =SFA g

(4)

where pKa is the dissociation constant for the acidic salt, SMS the solubility of the salt, and SFA the solubility of the free acid. As schematically represented in Figure 11 for the salt of a weak base,39 when the solution pH > pHmax the free base will be the equilibrium solid phase and disproportionation will be energetically favored, and when pH < pHmax the salt will be the equilibrium solid. Correspondingly, for the salt of a weak acid when the solution pH < pHmax the free acid will be the equilibrium solid, while at pHmax > pH, the salt will be the equilibrium solid. Commonly, the drop in solubility observed for both curves is due to the common ion effect. Many studies of various crystalline API salts have detected solidstate disproportionation when the API: (1) is stored at elevated temperatures and RH; (2) undergoes processing, for example, reduction of particle size and tablet compaction; and (3) is physically blended with various types of excipients, particularly those with intrinsic acidic or basic properties. Important pharmaceutically related effects of disproportionation on API solid-state properties include formation of volatile free acids or free bases, such as gaseous HCl, and low melting point organic amines capable of sublimation,42 and the formation of very water-insoluble free crystalline forms of the API which reduce dissolution.43 Of particular interest in the context of the current analysis is the essential role played by exposure of salts to various RHs, and to differences, which occur, with different counterions of the same salt. Most significantly, as might be expected, when salts are highly dried to remove measurable sorbed water, no solid-state disproportionation occurs. With regard to the effects of the amount of water vapor sorption at a particular RH, in the context of Equations 2 and 3, we would certainly expect measurable disproportionation for deliquescing salts under appropriate pH conditions, on account of the presence of the liquid water produced. Thus, in such cases we would readily expect the equilibrium solubility of various species, and the pH of a saturated solution to be useful in ascertaining where the pH of the system might be with regard to pHmax. What is

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Figure 11. Schematic illustration of the solubility versus pH for a weak base (left side) and a weak acid (right side) with the establishment of pHmax. Reproduced with permission from Thakral and Kelly.39

particularly interesting is the fact that measurable disproportionation also has been observed at amounts of adsorption well below the deliquescence point where very small amounts of water, at best, can exist as “liquid-like layers,” as described above for the alkali halides and various procaine salts.22,34 To illustrate these phenomena, we have chosen to follow one study, which comprehensively examined the solid-state disproportionation of the mesylate and napsylate salts of miconazole and benzocaine below the deliquescence point at 25 C and an RH of 57%.36 The expected pHmax for the various salts in solution was determined from Equation 2 using data presented in Table 4. Also included in Table 4 are values of the saturated solution pH of the various salts, believed to be an estimate of the expected microenvironmental pHmax in the sorbed water layers. Note that the measured pH for the saturated solutions of the various salts was either very close to or less than pHmax. Water vapor sorption isotherms for the 4 salts at 25 C indicate that more water sorption occurs with both mesylate salts than with the corresponding napsylate salts. From surface area measurements of the various salts, it was possible to calculate the number of “equivalent sorbed monolayers” occurring at 57% RH (below any deliquescence), and these are miconazole mesylate (9.6), miconazole napsylate (0.8), benzocaine mesylate (3.6), and benzocaine napsylate (0.6). Role of Excipients in Disproportionation Of particular interest is the importance of physically mixed excipients, which exhibit acidic or basic properties that can change the pH of the sorbed water relative to pHmax. Such excipients include basic materials such as magnesium oxide (a component of commercial magnesium stearate) and sodium croscarmellose, and acidic excipients such as citric and tartaric acids. A number of studies on the effects of such excipients indicate that water vapor sorbed into the bulk solid, for example, formation of crystal hydrates or absorption into the bulk amorphous structure is not a significant factor in solid-state disproportionation.36,38 Rather, it appears that only water located at the surface of the salts and excipients is involved in this process. To look at this more closely, we can choose one representative basic excipient used in the study of the miconazole and benzocaine salts, croscarmellose sodium (CrosNa), a widely used tablet disintegrant.36 The pH of 10% aqueous slurry of this very water-insoluble material is 6.85, and the water vapor sorption at 57% RH is 19.67% w/w, indicating relatively strong basicity and a high degree of polarity, respectively. Thus, with such a high “microenvironmental pH” relative to the pHmax of the various salts, it is not surprising that the mol% conversion of

salts to free base at a 50/50 w/w physical blend after 24 h is miconazole mesylate (25.7%), miconazole napsylate (no measurable reaction), benzocaine mesylate (21.2%), and benzocaine napsylate (no measurable reaction). Clearly, from the results, it is apparent that the more water-soluble mesylate salts, showing more significant sorption of water vapor at 57% RH than the napsylate salts undergo greater disproportionation in the presence of a physical blend with CrosNa. Because CrosNa is a highly amorphous solid, we would expect the large amount of water vapor sorption at 57% RH to be primarily due to the water, which is absorbed into the bulk phase with, of course, some level of limited surface adsorption of a few equivalent molecular layers. We can now ask 2 questions. How does the limited amount of sorption by the crystalline mesylate salts of miconazole and benzocaine, and the small amount of water expected to be at the surface of CrosNa, support the disproportionation of the mesylate salts? In particular, how does a physically blended excipient produce surface layers of sorbed water molecules that can “connect” with those of the API to produce a basic “surface solution,” capable of supporting the disproportionation reaction? As has been demonstrated above for the solid-state degradation of various procaine salts,34 it appears that sorbed layers of water at relatively low levels produce “liquid-like layers” capable of acting as a “solvent” for water-soluble species and supporting acid-base reactions. The exact structure of water in the sorbed layer and the extent to which high energy sites such as defects and local solid disorder (amorphous) influence this structure are not clear. For example, to what extent would preferential sorption by disordered surface regions plasticize the surface to create increased molecular diffusion along the surface? Whatever the exact mechanisms, worthy of more systematic study, there is no doubt from the many experimental observations reported in the literature that this limited amount of sorption of water can provide a vehicle for molecular diffusion and chemical reactivity. When physically mixing powdered material, barring any strong interaction between the solids at the point of contact, we would expect, from Equation 1, as discussed above, that at a particular RH the total amount of water sorbed at a given RH would be the weighted average of the amounts taken up by each component.19,20 For one component to “dissolve” and interact with the other component in sorbed water at constant RH requires that there must be physical contact between the surfaces of the 2 components, including any sorbed water at the surface. It is well established, for example, that a blend of 2 deliquescent solids, a and b, with RHoa

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and RHob will exhibit an RHoab below that of each component, described by the following equation44:

½RHoab =100 ¼ ½RHoa =100 þ ½RHob =100

(5)

Because this relationship arises from solution theory, we can conclude that contact between each of the components dissolved in the liquid water on the surface does occur. In the cases of chemical degradation and disproportionation of crystalline salts in the presence of sorbed water below deliquescence and various excipients, therefore, we would expect the surface area of contact between particles to be critical for the “mixing of the sorbed layers of each component.” Indeed, many studies have shown the importance of the ratio of the drug to excipient (more contact of excipient with API, the greater the effect), reduced particle size (hence, greater surface area of contact), and powder compaction (increased bulk density), all leading to enhanced surface-to surface contact in determining chemical reactivity and other drug-excipient interactions. Exactly how sorbed water associated with the API salt and the excipient makes contact sufficient to bring about reactivity is not clear. One possible mechanism for the drug-excipient contact and interaction might be the tendencies for local disordered regions of the drug and excipient to form a miscible molecular mixture with enhanced molecular mobility assisted by water sorption in these regions. Another contributing possibility might be the occurrence of capillary condensation at the points of contact, which produces bulk liquid water.6 Whatever the exact mechanism, we can conclude that water vapor sorbed below the point of deliquescence can act as a “liquid-like” medium, which can interact with various water-soluble salts, and acidic or basic excipients, to promote proton transfer as required of many acid-base reactions. Knowing the exact structure of such “liquid-like layers” of water sorbed by organic molecules, and exactly how it can bring about solid-state reactivity, would be immensely valuable for understanding the broader issues of solid-state drug-excipient interactions in the presence of low amounts of surface water. Water Vapor Sorption by Amorphous Salts Occurrence of Amorphous Salts Amorphous forms of pharmaceutical salts occur for a number of possible reasons. First is the case where stable anhydrous or hydrate forms are not obtainable through solution-mediated and solid-state crystallization. For example, quinapril HCl only is crystallizable as organic solvates, for example, acetonitrile or ethyl acetate, which rapidly convert to the amorphous form upon exposure to ambient conditions.45 Second, processing crystalline salts, for example, by milling, compaction, or drying, often produces residual amorphous structure in the solid with tendencies for increased instability (see discussion of process-induced disorder above).46 Third, acidic and basic APIs and excipients can form amorphous salts, in situ, when the API formulation undergoes lyophilization, melt extrusion. or spray drying.47

counterion in determining Tg are seen by comparing the Tg values for the amorphous Liþ, Naþ, Kþ, Rbþ, and Csþ salts of indomethacin, which are 140 C, 121 C, 109 C, 77 C, and 69 C, respectively.50 The order of Tg for these alkali metal salts is consistent with the order of a decreasing ionic radius of the counterion, and hence an increased charge density and greater electrostatic interaction energies. Similar effects of the various alkali metal counterions have been observed for the Tg’ of the freeze concentrates obtained during the lyophilization of the alkali metal salts of ganciclovir.51 A comprehensive study of propranolol and nicardipine salts, prepared with a variety of organic counterions, indicates that the Tg of these salts is much greater than their free bases, but still quite low relative to ambient temperatures. See, for example, in Table 5 values of Tg obtained for 10 salts of propranolol and its free base. Detailed multivariant analysis indicated good correlations with counterion pKa, and with its electrophilicity index.52 Effects of Plasticization As we might expect, water vapor sorption by amorphous salts is generally greater than that of the corresponding amorphous free form, as shown in Figure 12 for indomethacin and it sodium salt.53 This is also true for the water vapor sorption at 2 RH values by propranolol salts and its free base, as shown in Table 5.52 Clearly, for some of the salts the very large amount of water sorbed most likely would suggest deliquescence, or perhaps, crystal hydrate formation. Note that the lowest amount of water sorbed, other than that of the free base, occurs with the relatively nonpolar organic counterions, such as the oxalate, besylate, and tosylate ions, and the much greater sorption with the more polar mesylate salt. It is also interesting to note the significant difference in sorbed water for the more polar salts at 55% RH and 75% RH. In summary, as discussed above in the general overview, the sorption of water vapor by amorphous solids generally has a very significant effect on the thermodynamic and kinetic properties of the solid relative to the crystalline form. Of greatest importance are the plasticizing effects of water, which tend to increase diffusional molecular mobility. In view of this, we would expect that the presence of sorbed water in the bulk phase of the amorphous salt would produce major changes in thermodynamic and kinetic physical chemical properties relative to those of the crystalline salt, for example, increased chemical degradation,54 greater rates of crystallization,55 greater rates of salt disproportionation,56 and deliquescence at lower critical RHs (RHo) than that of the crystalline form.57 Water Vapor Sorption of Cocrystals Overview Cocrystals, as the name implies, represent multicomponent solid forms that exist in a single-phase crystalline lattice, made up Table 5 The Glass Transition Temperature (Tg), pKa, and Water Vapor Sorption (% w/w) at 55% RH and 75% RH for Various Salts of Propranolol and Its Free Base52 Salt

Glass Transition Temperature A key characteristic of amorphous salts that adds to the overall physical and chemical properties expected of the corresponding crystal is the enhanced diffusional molecular mobility that exists, and the major change in mobility that occurs at the glass transition temperature. Generally, the glass transition temperature of an amorphous salt is significantly greater than that of the corresponding free form. For example, the Tg of indomethacin is 42 C, while that of the sodium salt is 121 C,48 while quinapril HCl has a Tg of 91 C, and that of the free base is 51 C.49 The significant contributions of the interactions between the ionized species and its

Hydrochloride Tosylate Mesylate Besylate Oxalate Phosphate Tartrate Citrate Acetate Free base

Tg ( C)

40 17 29 e 32 11 23 21 9 9

pKa

6.0 1.3 1.2 0.7 1.3 2.0 3.0 3.1 7.6 e

Water Vapor Sorption % w/w 55% RH

75% RH

2.20 0.34 4.00 0.50 0.30 16.50 1.20 8.50 28.10 0.023

1.80 0.39 23.50 11.10 0.50 21.30 5.34 13.10 61.90 2.54

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Figure 12. Water vapor sorption at 30 C of amorphous sodium indomethacin (dark squares) and amorphous free acid form of indomethacin (dark circles). Reproduced with permission from Tong.53

of 2 or more components, usually in some distinct molar stoichiometry.58,59 Interest in forming pharmaceutical cocrystals comes from the possibility of increasing API solubility, dissolution, and oral bioavailability using water-soluble coformers,60,61 improving physical and chemical stability,62,63 and improving solid mechanical properties for pharmaceutical processing.64,65 A number of binary combinations of different therapeutic agents also have been prepared as “multidrug” cocrystals.66 Typically, cocrystals form by solution-mediated crystallization, liquid-assisted grinding, and neat grinding at elevated RHs. Unlike crystalline salts of weak acids and bases, where proton transfer between components must take place, cocrystals can form with any pair of molecules containing electron donors and acceptors, which can exist as a solid at ambient conditions and form thermodynamically stable hydrogen bonding motifs. Additional intermolecular interactions can occur with compounds containing aromatic rings capable of stacking and undergoing charge transfer interactions. As with all API crystals, there is the potential for cocrystals to exist in a number of polymorphic forms and hydrates and solvates. For example, a 1:1 cocrystal of caffeine and anthranilic acid was shown to exist in 2 polymorphic forms, 2 hydrates, and 7 solvates.67 Caffeine forms a nonstoichiometric (0.8:1) channel hydrate, such that the water in the channels tends to move relatively easily under the stress of changing RH or temperature.68 Similarly, the caffeine:anthranilic acid cocrystal exhibits “zigzag” channel-like structure, which likely allows the ready formation of the various hydrates and solvates. Polarity of the Coformer and Water Vapor Sorption Of particular interest in this article are the water vapor sorption characteristics of cocrystals and their impact on chemical and physical stability of the cocrystal compared to that of the crystalline API alone. From a review of the pharmaceutical cocrystal literature with regard to improving the bioavailability of poorly soluble drugs, a general pattern appears to exist, wherein, most often, cocrystals of a poorly water-soluble API are produced, crystal structures are determined, and the aqueous dissolution and bioavailability relative to that of the parent drug are measured. Furthermore, they generally are stored at various RHs for assessment of potential water-induced chemical and physical transformations. Such transformations can include crystal hydrate formation, chemical degradation, and cocrystal dissociation. Let us first consider the situation where the API, as a salt with adequate solubility and bioavailability, is stored at

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elevated RH and in situ forms an undesirable crystal hydrate or inadvertently transforms from one stoichiometry to another. Such a case, for example, is naproxen discussed above (Fig. 3), which, although having acceptable dissolution as the marketed sodium naproxen when compared to free naproxen, forms a variety of hydrates at different RH and temperature in situ.7 The potential problem here is whether one can adequately control RH to avoid the inadvertent transformation from the most stable tetrahydrate to other hydrate forms. One approach reported69 used an adequately water-soluble 2:1 cocrystal of naproxen-nicotinamide with a reduced ability to sorb water vapor relative to the naproxen salt. Figure 13 illustrates the relative dissolution rates of free naproxen, a 2:1 cocrystal of naproxen and nicotinamide, a 2:1 physical blend of naproxen and nicotinamide, where the dissolution is enhanced by forming the cocrystal. Figure 14 compares the water vapor sorption of sodium naproxen and that of the naproxen-nicotinamide cocrystal. Note the same behavior observed for sodium naproxen in the first sorption profile in Figure 3, and the very insignificant amount of water adsorbed by the cocrystal over the entire range of RH values in Figure 14. Thus, through cocrystal formation, dissolution enhancement of naproxen occurred, but now there was a lack of water vapor sorption and no inadvertent hydration and dehydration, as occurred with sodium naproxen. Other examples of where a cocrystal with enhanced aqueous dissolution did not sorb significant water vapor over a wide range of RH include the following: 1:1 indomethacinsaccharin,70 1:1 AMG 517-sorbic acid,61 and a 2:1 cocrystal of 2-[4(chloro-2-fluorophenoxy)phenyl]pyrimidine-4-carboxamide and glutaric acid.60 In the case of the naproxen cocrystal mentioned above and the other examples where water vapor sorption is very small, we can assume that the API itself, in each case, is fairly hydrophobic, and thus would not be expected to sorb much water beyond surface adsorption as a crystal. Consequently, it is the polarity of the coformer and the crystal structure of the cocrystal that likely are the major factors in maintaining low levels of sorption.

Dissociation and Physical Stability In some cases we can expect that the polarity of the coformer could lead to significant water vapor sorption, leading to the cocrystal hydrate, to deliquescence, or to the dissociation of the cocrystal to allow hydrate formation or deliquescence of the individual components. To illustrate such possibilities, let us first consider a comprehensive study by Jayasankar et al.71 of 2 theophylline-citric acid cocrystals, a 1:1 anhydrous form and a 1:1:1 cocrystal hydrate, prepared by crystallization from a suspension of individual components at various water activities, aw. Studies of the phase behavior revealed that both water and citric acid play a role in

Figure 13. Dissolution profiles for: (a) crystalline naproxen, (b) 1:1 physical mixture of naproxen and nicotinamide; and (c) 1:1 cocrystal of naproxen-nicotinamide. Reproduced with permission from Ando et al.69

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Figure 14. Water vapor sorption isotherms at 25 C for: (a) 1:1 naproxen-nicotinamide cocrystal and (b) crystalline sodium naproxen. Reproduced with permission from Ando et al.69

determining phase stability. The critical aw for the transformation of the anhydrous cocrystal to the hydrate was 0.8, that is, the anhydrous form is favored below a water activity of 0.8 and the hydrate above this value. Both cocrystal forms (anhydrous and hydrate) were exposed to 85% and 98% RH (recall aw is approximately RH/100) for 5 months at room temperature and examined for any phase changes. Both cocrystal forms were stable at 85% RH, but underwent dissociation at 98% RH to form the theophylline 1:1 monohydrate. Presumably, the different critical water activity with sorption than with the solution-mediated system was due to different kinetics for the transformation. Furthermore, water sorption at 85% RH was in the range of 4%-6% w/w, while at 98% RH the amount sorbed was >70% w/w, with the appearance of deliquescence. Similarly, a caffeine-citric acid cocrystal deliquesces at 98% RH and dissociates to form the caffeine hydrate.72 A further indication of the importance of coformer properties with regard to water vapor sorptioneinduced dissociation of cocrystals is found in the work of Trask et al.73 They prepared 5 anhydrous caffeine cocrystals with a series of dicarboxylic acids: (A) oxalic acid (2:1), (B) malonic acid (2:1), (C) maleic acid (1:1), and 2 cocrystal polymorphs (E) and (F) with glutaric acid (1:1) and tested their phase stability at a number of RHs (0%, 43%, 75%, and 98% RH) for up to 7 weeks. In Table 6 are listed the pKa, aqueous solubility, and critical deliquescence points (RHo) for each of the dicarboxylic acids. The following general trends of stability can be summarized as follows. Cocrystal A (oxalic acid) was found to be stable at all RH values and at all time periods. Cocrystal B (malonic acid) was stable up to 75% RH, showed dissociation and caffeine hydrate formation at 98% RH after 1 week. By 7 weeks, it had deliquesced. Cocrystal C (2:1 maleic acid) was stable up to 75% RH with the formation of caffeine hydrate at 98% RH at 1 day. The same results were obtained with the 1:1 maleic acid cocrystal (D) as was observed with the 2:1 cocrystal. Cocrystal E (1:1 glutaric acid) converted to its conformational polymorphic form, F, in 3 days at 43% RH. Cocrystal F remained unchanged at 75% RH but exhibited dissociation and caffeine hydrate formation with deliquescence at

Table 6 pKa1, pKa2, Aqueous Solubility, and Critical Deliquescence Point (RHo) for Various Coformer Dicarboxylic Acids Used to Form Caffeine Cocrystals73 Coformer

pKa1

pKa2

Aqueous Solubility (m)

RHo (%)

Oxalic acid Glutaric acid Malonic acid Maleic acid

1.2 2.8 1.9 4.3

4.2 5.7 6.1 5.6

1.3 10.7 15.3 19.5

98 88 73 56

98% RH. These results clearly indicate the least soluble acid, oxalic acid, showing the highest RHo and being the strongest acid was most resistant to water sorption leading to deliquescence, which occurred in the other cases. The propensity of Cocrystal E to convert to Cocrystal F in the presence of water is related to the different free energies exhibited by E and F, making F the more thermodynamically stable polymorph. The same instability patterns of forms B, C, D, and F at 98% RH would suggest that deliquescence at this RH was a major factor in bringing about dissociation and caffeine hydrate formation. In this study and the one presented above,71 consideration was given to the possibility that trace amounts of free acid were responsible for the initiation of deliquescence (Table 6). Another interesting example where cocrystal formation can influence the physical stability of an API is the drug, S-oxiracetam, a chiral compound available as the enantiomeric forms, S-ox and Rox, as well as the racemate, RS-ox.74 It was previously determined that S-ox performed therapeutically better than RS-ox; however, the S-ox crystal readily deliquesced at 87% RH, while RS-ox did not deliquesce until being exposed to 98% RH. Checking for the possible applicability of Wallach’s rule, which states that a racemic crystal generally has a higher density than its chiral counterpart, and hence should have greater molecular packing and physical stability,75,76 they prepared cocrystals of 2 chiral and 2 racemic 1:1 pairs of S-oxiracetam, using gallic acid (ga) and 3,4-dihydroxybenzoic acid (pa). They carried out detailed crystal structure determinations and exposed the samples to various RHs for up to 8 weeks. As expected, S-ox alone deliquesced at 87% RH, and RS-ox did not deliquesce until 3 days at 98% RH. However, the gallic acid cocrystals of S-ox and RS-ox were completely stable at 98% RH for up to 8 weeks. The RS-ox:pa cocrystal exhibited better stability than the S-ox:pa cocrystal, which showed deliquesce at 87% RH after 3 days, but still showed some deliquescence at 98% RH after 14 days. Measurements of water vapor sorption of gallic acid and 3,4dihydroxybenzoic acid indicated that both form hydrates at 87% and 98% RH, leading to the conclusion that dissociation to the coformer hydrates was not occurring. Estimating the densities of Sox and RS-ox cocrystals from the crystal structure indicated no difference between the S-ox:ga and RS-ox:ga cocrystals, but significantly greater density for the RS-ox:pa cocrystal than its Sox:pa counterpart. Therefore, it was concluded that Wallach’s rule appeared to be applicable to the 3,4-dihydroxybenzoic acid cocrystals, but not to the gallic acid cocrystals. Significant differences in hydrogen bonding patterns between the gallic acid and 3,4dihydroxybenzoic acid cocrystals were observed, but no clear picture could be determined about different responses of the 2 coformers at elevated RH. Chemical Stability Another potential use of the cocrystal form, not often reported, is to improve chemical stability while not negatively affecting the bioavailability. As an example of this strategy, let us consider a study of the drug temozolomide (TMZ) which is a prodrug of the active agent, 5-(3-monomethyl-1-triazeno) imidazole-4carboxamide (MITC).77 In solution, TMZ hydrolyzes to MITC at pH >7, which breaks down further to form a colored byproduct, 5aminoimidazole-4-carboxamide (AIC). When TMZ is prepared in a solid dosage form and exposed to 40 C/75% RH a color develops, indicative of solid-state degradation. To possibly offset this degradation, anhydrous cocrystals with 6 organic acids were prepared with the rational that an acid species in the crystal lattice would inhibit the base catalyzed reactions. The various acids included were oxalic acid (OA), salicylic acid (SAC), succinic acid (SA), malic acid (MA), anthranilic acid (ANA), and tartaric acid (TA). The dicarboxylic acids formed 2:1 API-acid cocrystals, while the

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monocarboxylic acids formed 1:1 API-acid cocrystals. Samples were stored at 40 C/75% RH and examined by X-ray powder diffraction (XRPD). TMZ alone started to form the TMZ hydrate after 2 weeks, while AIC appeared after 4-5 weeks. In contrast, TMZ:SA, TMZ:OX, TMZ:SAC, and TMZ:MA were stable for 28 weeks, while TMZ:ANA and TMZ:TA showed instability at 3 and 8 weeks, respectively. In all cases, the cocrystal formed involved a hydrogenbonded hetero-synthon between the TMZ carboxamide group and the carboxyl group of the acids. No indication was given of any differences in the water vapor sorbed by the cocrystals formed from different acids under these conditions, or of any cocrystal structural differences. It is apparent from the representative studies used in this analysis, therefore, that the properties of the coformer used to form soluble cocrystals, particularly more polar molecules, influence the chemical and physical stability of the API during exposure to high RH. Consequently, when using a polar coformer to enhance dissolution and bioavailability, it is important to weigh any beneficial effects against the possible detrimental effects caused by water vapor sorption, for example, crystal hydrates, deliquescence, and dissociation. Water Vapor Sorption by Coamorphous Systems Overview Coamorphous materials contain amorphous small molecule coformers which form a miscible mixture with the amorphous API.78 The goal is to increase solubility by making the API amorphous and subsequently to increase the stability by adding the small molecule coformer (physical and chemical). One explanation for the increased stability is that the coformer will increase relaxation times resulting in more stable systems.79 Coamorphous samples can be made with common methods that produce amorphous materials, such as spray drying, hot melt extrusion, and rapid precipitation. Many coamorphous samples are found during cocrystal screens but are not examined further because a crystalline system is desired. Further characterization of these amorphous screening samples could result in another solid form that could be developed. There are several examples where both coamorphous and cocrystal systems have been found (Table 7); however, it is not clear why one is produced over another and this is an area in need of further research. The list of coformers used for coamorphous and cocrystal systems is similar, including counterions and other APIs. As discussed for cocrystals, the polarity of the coformer and the hydrophobicity of the API can play a role in moisture uptake. The tendency of the polar small molecule coformers to crystallize from the amorphous state at elevated RH can be a significant disadvantage. For example, the water sorption of several amino acids has resulted in conversion to a crystal hydrate or deliquescence.91 It is expected that amino acids and other polar coformers could exhibit significant instabilities at lower RH due to the uptake of water and more work is needed to better understand how this relates to the coamorphous systems. Studies comparing the water sorption profiles of coamorphous systems containing different coformers were not found in the literature.

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Composition One possible advantage of coamorphous materials is the ability to exist in nonstoichiometric forms which allows a wide variety of compositions. This can be beneficial when a high drug load and low coformer content are needed for high dose formulations. Many coamorphous systems are produced in stoichiometric ratios; however, a few studies have used a wider range of coformer and API ratios.89,92-95 Water sorption studies of coamorphous systems with varying compositions have not been studied extensively. One study collected water isotherms on disodium cromoglycate with various amounts of L-lysine (2%, 5%, 10%, 20%, and 40%).96 All the isotherms showed significant uptake at 90% RH, ranging from ~30% to 50% water (while there was no mention of deliquescence in the report, the high water content might suggest that deliquescence has occurred). Samples containing the most lysine exhibited the lowest sorption values. Significant hysteresis was also observed indicating slower escape of water during desorption. XRPD analysis on samples stored at 75% RH for 24 h resulted in amorphous patterns, indicating that the samples were stable under these conditions. The amount of lysine in the sample was correlated to particle agglomeration of the disodium cromoglycate based on scanning electron microscopy analysis, with higher lysine contents showing less agglomeration of samples at 60%-75% RH. The studies suggested that 20% lysine successfully prevented the agglomeration of the API, resulting in improved aerosolization performance. Salbutamol sulfate and lysine coamorphous samples were prepared with compositions ranging from 0% to 100% API.97 Amorphous salbutamol sulfate showed recrystallization at 60% RH in the sorption isotherm, with an initial weight gain of about 10% water and subsequent drop to about 1% water, suggesting crystallization. The coamorphous materials show that samples with higher lysine contents exhibited lower amounts of sorbed water and the recrystallization moved to 70% RH, suggesting a more stable solid when compared to the amorphous API. Another study reported the water uptake of coamorphous paracetamol and citric acid (0%, 25%, 50%, 75%, and 100% paracetamol) at 43% RH over 18 weeks.92 As shown in Figure 15, the paracetamol sample exhibited the least water sorption and the citric acid sample exhibited up to 4% water. The 25% paracetamol sample sorbed the highest amount of water (8.2%), which was higher than either of the individual components. The 50% paracetamol sample exhibited a similar water uptake to the citric acid sample and the 75% RH sample exhibited the lowest water vapor sorption compared to the other coamorphous samples. These data indicated that different compositions impacted the amount of water sorbed and did not directly correlate to theoretical values based on the API/coformer ratios. For all samples, crystalline materials were observed under humid conditions and the water contents were reduced when crystalline material was present. It was reported that interactions between the paracetamol and citric acid were observed in the infrared spectra of the coamorphous samples and it was speculated that water may break or disrupt hydrogen bonds between the API and coformer. This allowed the molecules to reconfigure and favored the crystallization of the components. The

Table 7 Examples of Coamorphous and Cocrystal Systems API

Coformer/API

Cocrystal Reference

Acetaminophen (paracetamol) Azelnidipine Indomethacin

Citric acid Maleic acid Arginine Lysine Lidocaine Lactose tetrahydrate

80,81

82-84

85

86

87

47

87

47

88

89

90

90

Sodium naproxen

Coamorphous Reference

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9

Pecent Water Uptake

8 7 6 25% paracetamol 5

50% paracetamol

4 citric acid 3 75% paracetamol

2 1

paracetamol

0 1

4

8

12

18

Weeks at 43% RH Figure 15. Water vapor sorption of paracetamol:citric acid coamorphous samples. Reproduced with permission from Hoppu et al.92

water values showed that the amorphous blends with high citric acid content contained the most water, and that the crystallization appeared faster on the surface of the samples due to water adsorption and plasticization at the surface. It should also be noted that the initial Tg values of the coamorphous materials ranged from ~11 C to 26 C, which is below room temperature and may indicate physically unstable materials; any plasticization from water could decrease the stability even more. Plasticization As with any amorphous system, one goal is to obtain a Tg that is approximately 50 higher than ambient or the proposed storage temperature. Many coformers exhibit lower Tg temperatures than the polymers used in ASDs, resulting in lower Tg values for coamorphous systems when compared to ASDs.78 This plasticization of the API with the coformer was demonstrated with glibenclamide amino acid coamorphous samples.98 When water is sorbed by a coamorphous sample, the Tg will be reduced even further, resulting in values that may be subambient or close to ambient, and may result in physically unstable solids. Sorption of water can also lead to a separation of the 2 materials, resulting in a physical mixture of the 2 amorphous phases.99 A 1:1 simvastatin-lysine coamorphous system exhibited an initial Tg value of 33.2 C.98 This Tg value was similar to amorphous

simvastatin produced in the study (32.5 C) and FTIR showed no stabilizing interactions between the 2 components. Within a few days at elevated humidity (ambient/60% RH), the coamorphous sample changed into a yellow waxy material, and the Tg was measured as 9.3 C after 2 weeks. This suggests that the sorbed water had transformed the sample into a supercooled liquid under these conditions. The system remained amorphous for ~56 days. Upon crystallization, the individual components were observed by XRPD and FTIR, and no evidence of a cocrystal was found. Physical Stability An important consideration for any amorphous system is the ability of the solid to remain amorphous under a variety of conditions. Various properties of an amorphous system might contribute to the stability, such as Tg,100 interactions between the components,101,102 molecular mobility,103 lack of water sorption (less mobility for crystallization),104 and crystallizability of the individual components.105 Although other factors may also contribute to stability, these factors are commonly considered for most amorphous systems. Many stability studies are performed under accelerated conditions (such as 40 C/75% RH), while others are performed under dry conditions. A limited number of studies are available for coamorphous systems, and some of these are summarized in Table 8. This is

Table 8 Summary of Coamorphous Stability Studies System

Stability Conditions

Time

Result

Reference

1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 1:1 2:1 1:2 1:1 2:1 1:2

RT/10% RH 25 C/60% RH 40 C/75% RH 25 C/0% RH, 25 C/60% RH Dry conditions RT, 40 C/dry conditions RT, 40 C/dry conditions RT, dry conditions RT, 43% RH RT, silica gel Ball milled, 40 C silica gel 25 C, 75% RH 4 C, 25 C/P2O5 4 C, 25 C/P2O5 4 C, 25 C/P2O5 4 C, 25 C, 40 C/silica gel 4 C, 25 C, 40 C/silica gel 4 C, 25 C, 40 C/silica gel

28 d 60 d 3 mo 3 mo 10 mo 332 d 332 d 27 wk 4 wk 5 mo 5 mo 7d 21 d 21 d 21 d 30 d 30 d 30 d

Coamorphous Coamorphous Coamorphous Coamorphous Coamorphous Coamorphous Coamorphous Coamorphous Paracetamol, citric acid Coamorphous Lactose, sodium naproxen Sodium naproxen lactose tetrahydrate cocrystal Coamorphous Coamorphous (4 C), indomethacin (25 C) Naproxen Coamorphous (4 C, 25 C), ranitidine HCl (40 C) Coamorphous (4 C), indomethacin (25 C, 40 C) Ranitidine HCl Form 2

106

Sulfathiazole:tartaric acid Lurasidone HCl:saccharin Repaglinide:saccharin Loratadine:citric acid Indomethacin:lysine Naproxen:arginine Naproxen:tryptophan Paracetamol:citric acid Paracetamol:citric acid Sodium naproxen:lactose Sodium naproxen:lactose Sodium naproxen:lactose Indomethacin:naproxen Indomethacin:naproxen Indomethacin:naproxen Indomethacin:ranitidine HCl Indomethacin:ranitidine HCl Indomethacin:ranitidine HCl

RT, room temperature. Bold: cocrystal formation; italics: separation into individual crystalline components.

107 108 109 47 47 47 92 92 90 90 90 110 110 110 111 111 111

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not meant to be a comprehensive list, but a sampling of the studies found in the literature. In general, dry conditions favor maintaining the amorphous nature of the samples. In the case of 1:1 naproxen:arginine and 1:1 naproxen:tryptophan, the samples remained amorphous for almost a year.112 Some systems such as 1:1 lurasidone HCl:saccharin,107 1:1 repaglinide:saccharin,108 and 1:1 loratadine:citric acid109 showed good short-term stability (2-3 months) at elevated humidities (60%-75% RH). Other coamorphous samples such as 1:1 paracetamol:citric acid readily crystallize into the individual components upon exposure to 43% RH over 4 weeks92 Two studies investigated different stoichiometries (1:1, 1:2, 2:1) of the coamorphous systems.110,111 An indomethacin:naproxen system110 showed peak shifts in the FTIR spectra that indicated molecular interactions between the components, resulting in a heterodimer. The Tgs showed negative deviations from the GordonTaylor equation, with the 1:1 ratio showing the largest deviation and strongest interaction. Samples with 1:2 and 2:1 ratios showed recrystallization of the excess drug upon storage at elevated temperature, while the 1:1 ratio samples remained amorphous. An indomethacin:ranitidine HCl coamorphous system111 also resulted in FTIR spectra that indicated interaction between the carboxylic acid and benzoyl amide of the indomethacin with the aci-nitro of ranitidine HCl. The Tg values were found to be low (29.3 C, 32.5 C, and 34.3 C) and were in good agreement with the Gordon-Taylor equation. After 30 days of storage, the excess component was found to crystallize from the 1:2 and 2:1 samples. The 1:1 sample was more stable, but did crystallize to ranitidine HCl after 30 days at 40 C. Although these studies did not specifically involve water sorption, they are good examples of the type of analyses that can be performed to better understand coamorphous systems. As seen in Table 8, one coamorphous system transformed into the cocrystal.90 Sodium naproxen:lactose coamorphous samples were produced by ball milling and spray drying, as well as dehydration of a sodium naproxen:lactose tetrahydrate cocrystal. Coamorphous samples stored at 25 C/75% RH quickly recrystallized into the cocrystal. This suggests that water is a necessary component for crystallization of the cocrystal and a determination of the critical RH would be key for determining storage and handling conditions for the coamorphous system. The ball milled sample stored under dry conditions at 40 C showed that signs of recrystallization into the individual

= API

= polymer 1

crystalline components composed mainly of crystalline lactose and minor peaks for sodium naproxen. Other samples stored under dry conditions for 5 months were found to remain amorphous. This system shows how processing (ball milled vs. spray drying vs. dehydration) and storage conditions (low and high RH) can produce a variety of solid-state transformations that must be understood when further developing these systems. Processing of components can also result in form changes. Cryogenic grinding of 1:1 carbamzepine:saccharin resulted in a mostly coamorphous material exhibiting a Tg of ~41 C. Upon exposure to RH (75% RH, room temperature), the sample transformed to the 1:1 cocrystal after 1 day.113 Ball milling the components at room temperature produced the cocrystal directly without a coamorphous intermediate. More work is needed on coamorphous materials to better understand the mechanisms for recrystallization into cocrystals or individual components during short and long-term exposure to water. These limited studies show that a variety of factors (such as Tg, interactions between components, miscibility, molecular mobility, processing conditions, storage conditions) with and without water need to be investigated to help determine the best conditions to produce and maintain stable coamorphous systems for development. Water Vapor Sorption by Amorphous Solid Dispersions Overview ASDs contain an amorphous API in a miscible solid mixture with a second component, commonly a polymer.114,115 The goal of the polymer is to help improve the physical stability of the amorphous solid (prevent solid-state API crystallization) and maintain supersaturation (prevent precipitation) upon administration. Because the polymers used in ASDs are usually water soluble and the APIs are quite hydrophobic, most water sorption observed in these multicomponent systems is due to the presence of the polymer. The water uptake of polymers can vary considerably, with polyvinylpyrrolidone (PVP) showing high water sorption and hypromellose acetate succinate (HPMCAS) showing low water sorption over most of the RH range used during processing.116 Other components have been used in ASDs, such as surfactants or a mixture of polymers and surfactants, creating ternary or even quaternary systems (Fig. 16). The addition of cyclodextrins or phospholipids

= polymer 2

= surfactant

Polymers and Surfactants

Water and solvent

Other Solubilization Mechanisms

Binary

Binary

= cyclodextrin

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Ternary

Binary

Ternary

Ternary

Quaternary

Ternary

= phospholipid

Figure 16. Schematic of possible amorphous solid dispersion systems. Reproduced with permission from Newman.117

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results in different solubilization mechanisms when compared to the polymers and surfactants, yet these systems are commonly included in ASD nomenclature. Dispersions are produced with common methods, such as spray drying, hot melt extrusion, and rapid precipitation. As with other amorphous systems, the increase in physical stability can be related to increased Tg values (antiplasticization), strong interactions between components, or by a dilution effect.118 Polymer concentrations of 70%-80% are commonly used during development, but lower polymer amounts can also provide protection, as shown, for an indomethacin:PVP dispersion containing 5% polymer that remained amorphous for up to about 5 month at 30 C and under dry conditions.119 Solution stability has been attributed to interactions between API and polymer in solution and the formation of a colloidal API-rich liquid phase that resists crystallization.116,120 Isotherm Prediction As shown in Equation 1, water sorption in a physical mixture of 2 components can be predicted based on the weight of water sorbed and the fraction of each component. Studies with miscible PVP dispersions and hydrophobic drugs (indomethacin, ursodeoxycholic acid, and indapamide), however, showed that the experimental isotherms were reduced relative to the predicted isotherms.121 These results indicated that formation of the dispersion altered the water sorption properties of the individual components. The greatest deviation was found close to the 1:1 drug:PVP monomer ratio, suggesting that the intermolecular interactions in ASDs significantly impacted the water uptake in these systems. A 3-component FloryHuggins model was used to successfully predict the isotherms for certain IMC:PVP compositions, as well as ASD systems containing sucrose and trehalose with PVP.122 Plasticization As discussed with other amorphous systems, water (or other components such as polymers or surfactants) can act as plasticizers, reducing the Tg of the system, increasing the molecular mobility, and possibly decreasing physical stability. For example, a dispersion containing a new chemical entity (SAR) and hypromellose phthalate (HPMCP) was exposed to various RH conditions, and the water

content and Tg values were measured.123 A single Tg was observed for all data points, indicating that a miscible system was maintained during the study. The dry dispersion exhibited a Tg of ~118 C, while a sample containing ~5.5% water resulted in a significantly lower Tg of ~63 C. When the Tg was plotted as a function of water content, a straight line was obtained which fit the theoretical Gordon-Taylor equation.124 In another study, solid-state nuclear magnetic resonance (SSNMR) spectroscopy was used to measure the relaxation times of sucrose-PVP dispersion and individual components before and after exposure to water.125 PVP alone showed an increase in mobility with an increase in RH; however, the polymer side chain motions in the ASD did not change with increasing RH. A similar result was found with sucrose, where the sucrose mobility in the ASD did not increase with RH even though amorphous sucrose alone did show a significant increase. The change in molecular mobility was attributed to hydrogen bonding between the components, which stabilized the dispersion even in the presence of water. Separation Into Amorphous Physical Mixtures Miscible amorphous dispersions, generally, will display distinct interactions between the components and exhibit only one Tg. Upon exposure to water, however, sometimes miscible ASDs can separate into physical mixtures of individual amorphous components (Fig. 17).104,126 Such physical mixtures of amorphous materials do not improve the physical stability, whereas most miscible systems show improved stability.126,127 Phase separation kinetics have also been measured for miscible ASD systems containing griseofulvin:PVP and indoprofen:PVP exposed to water vapor.128 Hydrophilicity and hydrophobicity of the polymer and drug have been shown to influence separation of miscible dispersions at elevated RH. One study involved dispersions containing a hydrophobic drug and a hydrophilic polymer.126 When water was added to the miscible dispersion, a ternary system containing drug, polymer, and water was produced. Apparently, the added water created a thermodynamically unfavorable environment, leading to separation into a drug-rich and polymer-rich region (path iii in Fig. 17). The favorable interactions between water and the hydrophilic polymer form a “cosolvent” system, and, as the water is increased, the miscibility of the hydrophobic amorphous drug and

Figure 17. Schematic representation of amorphous phase separation and crystallization from amorphous solid dispersions. (i) A solid dispersion is exposed to moisture and can (ii) crystallize immediately from the one-phase ternary (water-drug-polymer) system or (iii) separate into individual amorphous phases resulting in drug-rich and polymer-rich regions, which ultimately produce (iv) crystalline drug. Reproduced with permission from Rumondor et al.126

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the polymer is decreased, leading to phase separation. When phase separation occurs, the inhibitory influence of the polymer is reduced, and crystallization can occur more readily. An example is a miscible dispersion containing the hydrophobic drug pimozide and the hydrophilic polymer PVP, where phase separation was observed at 54% RH. In contrast, indomethacin:PVP and ketoprofen:PVP miscible dispersions were stable for weeks at high RH (94% RH) and did not show signs of phase separation. It was found that the phase separation was related to the strength of the drug-polymer interactions; model drugs with NH moieties (nifedipine, felodipine, droperidol, and pimozide) were more sensitive to moistureinduced immiscibility than those containing COOH functions (ketoprofen and indomethacin). The stronger bonds between the carbonyl groups of the drug and polymer made the dispersion less susceptible to bond breaking when water was introduced into the system, while the weaker NH and COOH bonding broke more readily in the presence of water. A second study with poorly watersoluble drugs reported that properties such as a less hydrophobic API, stronger drug-polymer interactions, and low dispersion water uptake reduce moisture-induced phase separation.104 It was shown, further, that the balance between thermodynamic factors, such as the enthalpy and entropy of mixing, in a ternary waterdrug-polymer system was dependent on the drug-polymer interactions and hygroscopicity of the polymer. Dispersions exhibiting strong drug-polymer interactions and a less hygroscopic polymer were less susceptible to phase separation, while more hydrophobic drugs were more susceptible to separation, even at low values of sorbed water.129,130 A low crystallization tendency105 of an API may also play a role in phase separation. BMS-817399 is a lipophilic slow crystallizer that is prepared as an ASD with PVP.131 Drug-polymer interactions were confirmed at low RH conditions (75% RH), but at high RH the bonding between the components was not evident and phase separation was observed. Two roles of water during moistureinduced phase separation were proposed: (1) disruption of the drug-polymer interactions and (2) plasticizing of the ASD to increase the molecular mobility and accelerate the amorphous phase separation. The surface of the solid exhibited enrichment of the hydrophobic drug after exposure to 95% RH. The low crystallinity potential of BMS-817399 helped retain the amorphous nature of the drug upon separation, and crystallization was not observed for the system. Crystallization of API Although many ASDs are physically stable and have been developed into marketed products,132 investigating possible crystallization is a critical step during development. Crystallization from the amorphous state is a combination of thermodynamic (free energy) and kinetic (molecular mobility) factors.133 As shown in Figure 17, 2 mechanisms have been proposed for API crystallization in ASDs due to water. The first, as discussed in the previous section, starts with water sorption and results in phase separation of the individual amorphous components. Once the amorphous API is a separate phase, the physical stability will commonly decrease resulting in crystallization (path i-iii-iv in Fig. 17). The other proposed mechanism is API crystallization directly from the ternary amorphous system containing API, polymer, and water (path i-ii-iv in Fig. 17). Both paths need to be avoided and studies are needed to understand possible API crystallization and determine the best methods to prevent it in drug products. The form that crystallizes from the ASD also needs to be determined, because both anhydrates and hydrates can be formed at particular elevated RH conditions. It is possible that different anhydrous and hydrated forms can be crystallized under various conditions, and this can be illustrated with ritonavir dispersions.134 Amorphous ritanovir samples

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stressed at elevated RH conditions crystallized into Form I, which was the kinetically favored metastable form. A ritonavir:HPMCAS dispersion equilibrated under the same conditions crystallized as a mixture of Form I and Form II, with the majority being Form II, the most stable form. Indomethacin:PVP dispersions also showed different indomethacin forms crystallizing from the dispersions and amorphous API. Pure amorphous indomethacin crystallized as the g form or mixtures of a and g forms, whereas the dispersions resulted in only the g form.135 These studies show that the forms produced from crystallization of the pure amorphous API did not provide the necessary information on subsequent crystallization from the ASD due to the possible effects of the polymer on the crystallization kinetics (diffusion barrier for nucleation and growth) or the impact of API configuration or hydrogen bonding with the polymer on the crystallization (interfacial energy). The method of preparation for the ASD can influence stability and crystallization. Dispersions of Compound X and polyvinylpyrrolidone vinyl acetate were prepared by spray drying and melt extrusion.136 Both dispersions were found to be miscible and had similar physical properties (Tg, compactibility, tabletability, disintegration, and dissolution). Differences were noted in surface area, morphological structure, powder densities, and flow characteristics. The higher surface area of the spray-dried dispersion (22 higher than the extruded ASD) resulted in a higher amount of sorbed water value (0.002 mol/g) compared to the extruded dispersion (0.00097 mol/g). Accelerated stability studies showed crystalline material in the spray-dried material after 1 day at 40 C/ 75% RH and 3 months at 50 C/51% RH (both open containers), whereas the extruded sample showed no crystallization under these conditions. The poor physical stability of the spray-dried dispersion was likely due the differences in surface area, with the higher amount of sorbed water measured for the spray-dried sample resulting in better plasticization of the system and faster crystallization. Another example of processing parameters affecting water sorption and stability involves a ternary griseofulvin:poly[N(2-hydroxypropyl) methacrylate] (PHPMA):PVP dispersion produced by spray drying in acetone/water versus acetone/methanol.137 Significantly different Tg values were measured for the ASDs, with the acetone/water preparation exhibiting a Tg of 83 C and the acetone/methanol showing a Tg of 103 C. The dispersion prepared from acetone/water sorbed more water (3%) compared to the acetone/methanol preparation (2.5%). Samples stored at ambient conditions for 13 weeks showed that the acetone:methanol sample was less stable and started to crystallize, which was unexpected based on the higher Tg and lower water sorption. The stability was explained by conformational variations of the polymer in the spray drying solutions which were then retained in the spray-dried solids. Stronger interactions between the PHPMA and PVP in the acetone:methanol system resulted in weaker interactions between the API and PHPMA, resulting in crystallization of griseofulvin. Water vapor sorption studies on formulated products, such as tablets, are also important during development to determine the effect of excipients on the crystallization potential over time. The physical stability of vemurafenib:HPMCAS dispersion was investigated for a formulated tablet under 3 stability conditions (25 C/60% RH, 30 C/75% RH, and 40 C/75% RH) with both open and closed conditions over 15 months.138 Water content and percent crystallinity were measured for all samples. Open conditions resulted in various amounts of crystallization ranging from ~2% crystalline API (25 C/60% RH) to 24% crystalline API (40 C/75% RH), with water contents ranging from 3.5% to 4.5%. The closed conditions resulted in moisture contents of 1.3%-1.6% water and crystalline contents <2% for 2 conditions (25 C/60% RH, 30 C/75% RH) and ~2% for 40 C/ 75% RH after 15 months. A combination of high humidity and

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temperature was found to destabilize this dispersion, and the control of moisture and temperature using protective packaging was essential to maintaining the amorphous nature of the API in the formulation. Specific protective packaging has been used for other systems to maintain amorphous API in ASDs.139 It should be noted that dissolution at 30 min was also correlated to crystalline API content produced during testing, but the bioavailability of vemurafenib:HPMCAS capsule formulation was found to be significantly higher for the amorphous dispersion (4.4-4.7 times) when compared to a micronized crystalline API capsule formulation, indicating that form changes in aqueous biological fluids was not a significant factor for performance of the ASD. Stability Prediction Stability prediction for long-term storage based on short-term stability or other data would help shorten the selection process for dispersions. A study with ritonavir:HPMCAS used isothermal calorimetry to monitor crystallization under various stress conditions, including elevated RH.134 An expanded Arrhenius model was employed, and a nonlinear approach was applied to minimize the propagation of errors associated with the estimates. The physical stability of this ASD was mainly governed by primary nucleation. The impact of polymers and moisture on the crystallization was modeled successfully, but further work was needed to understand the impact of temperature and water on nucleation and crystallization to improve the model. In another study, API (SAR):HPMCP dispersions were exposed to various temperature and humidity conditions below and above the Tg until crystallization of the API was observed.123 The data collected above Tg were then extrapolated below Tg, resulting in a qualitative trend. When analyzing the data further, it was found that the temperature for the onset of crystallization varied linearly with the Tg/T ratio. Statistical analysis showed that the data obtained at the highest temperature/humidity conditions (crystallization observed in less than 3 months) could be extrapolated over 15 months. These data were used to design appropriate temperature and humidity conditions for longterm storage to prevent crystallization. An efavirenz-PVP dispersion was used to develop a kinetic model capable of predicting the physical stability.140 Recrystallization kinetics were measured for the initial ASDs stored at controlled temperature and RH, and the data were fitted using a new kinetic model to estimate the recrystallization rate constant. The recrystallization rate constant was found to increase linearly with RH, and polymer content inhibited the recrystallization process by increasing the crystallization activation energy and decreasing crystallinity. The model was validated with experimental data and showed accurate predictions of stability. Although these 3 examples show the utility of stability prediction, additional work is needed in this area to determine if a more global approach could be applied to a larger number of dispersion systems. Summary This examination of water vapor sorption by multicomponent solids of pharmaceutical interest has shown that crystalline salts; cocrystals; and amorphous salts, coamorphous mixtures, and ASDs are subject to the same general effects of sorbed water as with any solid by means of adsorption, absorption, crystal anhydrate/hydrate transformations, deliquescence, and capillary condensation. What makes multicomponent solids unique in this regard, however, is that they consist of an API and a second component, that is, counterions, or small molecule or polymer coformers, where the properties of the second component greatly influence the ultimate behavior of the solid in the presence of water vapor. As a review of pertinent literature has revealed, these multicomponent systems

are prepared primarily because a crystalline API alone lacks adequate aqueous solubility, dissolution, and bioavailability. To this end, then, although the counterion or coformer should have significant polarity to enhance dissolution, this enhanced polarity will contribute to significant water vapor sorption, which in turn can lead to physical and chemical instabilities. Such instabilities include disproportionation of salts, and dissociation of cocrystals, coamorphous mixtures, and ASDs. It appears clear from this examination that water adsorption of crystalline multicomponent solids, below the deliquescence point, is capable of initiating chemical and physical instabilities at levels of water equivalent to just a few molecular layers in what have been referred to as “liquid-like layers.” It is also apparent that this adsorbed water can support pH effects on stability, as well as enhance drug-excipient interactions. With regard to amorphous salts, coamorphous mixtures, and ASDs, we see the importance of sorbed water as a plasticizer in bringing about increased molecular mobility and the possibility of subsequent phase separation and crystallization. In preparing multicomponent solids, therefore, it is important to measure the water vapor sorption isotherm of the counterion or coformer, to understand the mode by which water is sorbed, and to anticipate and correct the possible instabilities that might occur.

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