Solubility crystal hydrates

A thorough understanding of the hydration profile for a solid forming a crystal hydrate is important for several reasons. First, since an anhydrate and hydrate(s) are distinct thermodynamic species, they will have different physical-chemical properties (e.g., solubility) that may affect bioavailability. Second, a desired hydrate species can be formed and used (and retained) simply by controlling the desired, established environmental conditions. Third, since significant quantities of water can be sorbed/liberated as a hydrate becomes hydrated/dehydrated, the physical-chemical properties of the immediate system (including other nearby solids) can be markedly affected. 

Blue leaflets turns pink in moist air hygroscopic the dihydrate is violet blue crystal the hexahydrate is pink monochnic crystal density 3.36, 2.48 and 1.92 g/cm for anhydrous salt, dihydrate and hexahydrate, respectively anhydrous salt melts at 740°C and vaporizes at 1,049°C vapor pressure 60 torr at 801°C the hexahydrate decomposes at 87°C the anhydrous salt and the hydrates are aU soluble in water, ethanol, acetone, and ether the solubility of hydrates in water is greater than the anhydrous salt. 

The importance of polymorphism in pharmaceuticals cannot be overemphasized. Some crystal structures contain molecules of water or solvents, known as hydrates or solvates, respectively, and they are also called as pseudopolymorphs. Identifying all relevant polymorphs and solvates at an early stage of development for new chemical entities has become a well-accepted concept in pharmaceutical industry. For poorly soluble compounds, understanding their polymorphic behavior is even more important since solubility, crystal shape, dissolution rate, and bioavailability may vary with the polymorphic form. Conversion of a drug substance to a more thermodynamically stable form in the formulation can signiLcantly increase the development cost or even result in product failure. 

Figure 14-6 A graph that illustrates the effect of temperature on the solubilities of some salts. Some compounds exist either as nonhydrated crystalline substances or as hydrated crystals. Hydrated and nonhydrated crystal forms of the same compounds often have different solubilities because of the different total forces of attraction in the solids. The discontinuities in the solubility curves for CaCl2 and Na2S04 are due to transitions between hydrated and nonhydrated crystal forms.

In general, polymorphs of a given compound have different physicochemical properties, such as melting point, solubility and density, so that the occurrence of polymorphism has important formulation, biopharmaceutical and chemical process implications. In addition to polymorphs, solvates (inclusion of the solvent of crystallization), hydrates (inclusion of water of crystallization) and amorphous forms (where no long-range order exists) may also exist. Figure 3.8 shows an example of the polymorphism of estrone (Busetta et al. 1973). 

Extensive lattice relaxation such that the surface is theorized to be highly disordered is the basis for explaining various physical chemical phenomena for organic small molecules. Water vapor sorption to crystal hydration or potentially to the point of deliquescence, physical instability, chemical reactivity, changes in solubility, and interparticle bonding during compaction (Kontny 1995 Nystrom 1996 Hancock 1997 Buckton 1999) are all dependent on the physical chemical characteristics of the morphologically important crystal surfaces. 

It has been noted from the earliest dissolution work [39] that, for many substances, the dissolution rate of an anhydrous phase usually exceeds that of any corresponding hydrate phase. These observations were explained by thermodynamics, where it was reasoned that the drug in the hydrates possessed a lower activity and would be in a more stable state relative to their anhydrous forms [74]. This general rule was found to hold for the previously discussed anhydrate/hydrate phases of theophylline [42,44,46], ampicillin [38], metronidazone benzoate [37], carbamazepine [34,36], glutethimide [75], and oxyphenbutazone [72], as well as for many other systems not mentioned here. In addition, among the hydrates of urapidil, the solubility decreases with increasing crystal hydration [58]. 

The solubility and hydration number of alkali feed, diffusion coefficient of reacted ions and interaction of ion pair affect the reactive crystallization. As a result of investigation of a number of relationships between these properties and crystallization characteristics, the key property of alkali feeds may be considered as the following. 

However, Kondo [3 ] is of the opinion that the hydrolysis of the glass in water occurs and the calcium ions are released initially to the liquid phase. Simultaneously on the surface of slag grains an acid, colloidal shell of silica-alumina gel is formed. This shell has low permeability and hence the further slag reaction with water is hindered. In the presence of Ca(OH)2, added as alkaline activator, the siUcon and aluminum from the shell are released to the solution (Fig. 8.2). The solubiUty of aluminum compounds becomes considerably increased in the solution of pH higher than 12.5 because in this condition the Al(OH) ions are formed [4]. The concentration of aluminum in the liquid phase is increasing because the calcium aluminates crystallize a httle later, primarily the C-S-H (1) is formed. Simultaneously the solubility of hydrates formed in this condition is reduced. 

The Krafft point is the temperature at which the solubility of hydrated surfactant crystals increases sharply with increasing temperature and forming micelles. This increase is so sharp that the solid hydrate dissolution temperature is essentially independent of concentration above the critical micelle concentration (cmc) and is therefore often called the Krafft point without specifying the surfactant concentration. The steep increase in solubility above the sharp bend is caused by micelle formation. Micelles exist only at the temperature designated as the Krafft point. This is a triple point at which surfactant mole-... 

Figure 11.14 Phase diagrams for theophylline-citric acid (THP-CTA) anhydrous and co-crystal hydrate in water at 25 °C. Eutectic points are indicated by (THP hydrate/THP-CTA hydrate), Ej (THP-CTA hydrate/THP-CTA anhydrous) and Ej, (THP-CTA anhydrous and CTA hydrate). Solubilities of THP hydrate and CTA hydrate are indicated by a and b in each plot. (a) Phase solubility diagram generated from measured eutectic points and models that describe co-crystal solubility behavior, (b) Schematic triangular phase diagram showing the stability domains for anhydrous and hydrated co-crystals with co-formers that modulate the water activity. Stability regions for the crystalline phases are 1, crystalline drug hydrate 2, co-crystal hydrate 3 anhydrous co-crystal 4, co-former hydrate 5, crystalline drug hydrate/co-crystal hydrate 6, anhydrous/hydrated cocrystals 7, anhydrous co-crystal/hydrated co-former.

At low and medium supersaturations the number of particles formed depends on the number of heterogeneous nuclei and usually does not exceed 10 cm (Fig. 1). Once formed, crystals enlarge by deposition of solute ions at the surface, surface diffusion to a suitable site, and incorporation into the crystal lattice (see also Sections II.A and II.B). Under these conditions strongly hydrated cations are likely to form different crystal hydrates with different solubilities. Crystals thus formed are coarser and contain less—primarily crystalline—water than... 

Very often, the phase in equilibrium with the liquid solution below the Krafft eutectic is a crystal hydrate instead of the dry crystal [49]. Although cryst hydrates are very significant with respect to many aspects of phase behavior, thermodynamic analysis, and engineering processes, their existence does not alter the form of the crystal solubility boundary. Crystal hydrates may introduce complexity of other kinds into the physic behavior of surfactants, however, especially if the equilibrium crystal hydrate is slow to form [48]. 

At a temperature just above the Krafft eutectic (Fig. 3), the solubility boundary that is encountered is almost invariably a liquid-crystal solubility boundary. (The only circumstance when this is not true is when a liquid-liquid miscibility gap intrudes, in which case a liquid solubility boundary is found [72].) In highly soluble surfactants, the span of the miscibility gap beyond a liquid-crystal solubility boundary is typically very small ( a few percent). The miscibility gap at the Krafft eutectic may span three-quarters of the diagram, and in the vertic part of this boundary the miscibility gap is nearly 100% (unless crystal hydrates intervene). In a typical anionic surfactant such as SDS, for example, the tie-line lengths in the low-temperature region are about 88.9%. Just below the ICrafft eutectic they are 50%, ut just above this eutectic they are about 2.5% [42]. 

Chemical properties. Phalloidin, in the form of fine colorless needles crystallizes, hydrated with 5 molecules of water. The crystals obtained from methanol are also fine needles, which melt with decomposition at 255-258° C. Phalloidin is soluble to 0.5% in water at 0° C. It is distinctly more soluble in hot water, very soluble in aqueous or absolute methanol, ethanol, butanol, and pyridine, and it is insoluble in other organic solvents. Phalloidin is dextrorotatory [a] = +55 7 (in methanol) (628). 

Zincill) chloride. ZnCl2, is the only important halide—it is prepared by standard methods, but cannot be obtained directly by heating the hydrated salt. It has a crystal lattice in which each zinc is surrounded tetrahedrally by four chloride ions, but the low melting point and solubility in organic solvents indicate some covalent... 

Properties. A suimnaiy of the chemical and physical properties of alkah-metal and ammonium fLuoroborates is given in Tables 2 and 3. Chemically these compounds differ from the transition-metal fLuoroborates usually separating in anhydrous form. This group is very soluble in water, except for the K, Rb, and Cs salts which ate only slighdy soluble. Many of the soluble salts crystallize as hydrates. 

X 10 J/T (5.71 //g) at room temperature. It is air stable at 25°C, but is slowly converted to Fe202 and bromine at 310°C. The light yellow to brown hydroscopic sohd is soluble ia water, alcohol, ether, and acetonitrile. Iron(II) bromide forms adducts with a wide range of donor molecules. Pale green nona-, hexa-, tetra-, and dihydrate species can be crystallized from aqueous solutions at different temperatures. A hydrate of variable water content,... 

Iron(III) bromide [10031-26-2], FeBr, is obtained by reaction of iron or inon(II) bromide with bromine at 170—200°C. The material is purified by sublimation ia a bromine atmosphere. The stmcture of inoa(III) bromide is analogous to that of inon(III) chloride. FeBr is less stable thermally than FeCl, as would be expected from the observation that Br is a stronger reductant than CF. Dissociation to inon(II) bromide and bromine is complete at ca 200°C. The hygroscopic, dark red, rhombic crystals of inon(III) bromide are readily soluble ia water, alcohol, ether, and acetic acid and are slightly soluble ia Hquid ammonia. Several hydrated species and a large number of adducts are known. Solutions of inon(III) bromide decompose to inon(II) bromide and bromine on boiling. Iron(III) bromide is used as a catalyst for the bromination of aromatic compounds. 

Mercurous Nitrate. Mercurous nitrate [10415-75-5] Hg2N20 or Hg2(N02)2, is a white monoclinic crystalline compound that is not very soluble in water but hydrolyzes to form a basic, yellow hydrate. This material is, however, soluble in cold, dilute nitric acid, and a solution is used as starting material for other water-insoluble mercurous salts. Mercurous nitrate is difficult to obtain in the pure state directly because some mercuric nitrate formation is almost unavoidable. When mercury is dissolved in hot dilute nitric acid, technical mercurous nitrate crystallizes on cooling. The use of excess mercury is helpful in reducing mercuric content, but an additional separation step is necessary. More concentrated nitric acid solutions should be avoided because these oxidize the mercurous to mercuric salt. Reagent-grade material is obtained by recrystaUization from dilute nitric acid in the presence of excess mercury. 

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