Ch2oh

Structure XV Sorbitol

hydroxyl groups cannot bond onto the water 'lattice' without causing it to distort considerably. This may be one explanation of the difference, although differences in the lattice energies of the crystals may also contribute.

Hydration of ionic species: water structure breakers and structure makers

The study of ionic solvation is complicated but is relevant in pharmaceutics because of the effect ions have on the solubility of other species. The forces between cations and water molecules are so strong that the cations may retain a layer of water molecules in their crystals. The effect of ions on water structure is complex and variable. All ions in water possess a layer of tightly bound water - the water molecules being directionally orientated. Four water molecules are in the bound layer of most monovalent, monatomic ions. The firmly held layer can be regarded as being in a 'frozen' condition around a positive ion. The water molecules could be orientated with all the hydrogen atoms of the water molecules pointing outwards (see Fig. 5.2). Because of this and because their orientation depends on the ion size, they cannot all participate in the normal tetrahedral arrangements of bulk water (see Section 6.3.1). For this to be feasible, two of the water molecules must be orientated with the hydrogens of the water molecules pointing in towards the ion. Inevitably, then, with cations and many small anions there tends to be a layer of water around the bound layer which is less ordered than bulk water (Fig. 5.2). Such ions, which include all the alkali and halide ions except Li+ and F0, are called

Alternative orientation to suit ion

In-between water molecule

Bulk water

Figure 5.2 Schematic diagram to indicate that, in the (hatched) region between the primary solvated ion and bulk water, the orientation of the 'in-between' water molecules must be a compromise between that which suits the ion (oxygen-facing ion) and that which suits the bulk water (hydrogen-facing ion).

Reproduced from J. O'M. Bockris and A. K. N. Reddy, Modern Electrochemistry, vol. 1, MacDonald, London, 1970, with permission.

Possible orientation to suit bulk water

Alternative orientation to suit ion

In-between water molecule

Bulk water

Figure 5.2 Schematic diagram to indicate that, in the (hatched) region between the primary solvated ion and bulk water, the orientation of the 'in-between' water molecules must be a compromise between that which suits the ion (oxygen-facing ion) and that which suits the bulk water (hydrogen-facing ion).

Reproduced from J. O'M. Bockris and A. K. N. Reddy, Modern Electrochemistry, vol. 1, MacDonald, London, 1970, with permission.

structure breakers. The size of the ion is important, as the surface area of the ion determines the constraints on the polarised water molecules. Many polyvalent ions, for example Al3+, increase the structured nature of water beyond the immediate hydration layer, and are therefore structure makers.

such as iodide, caesium and tetraalkylammo-nium ions. The solvation numbers decrease with increase of ionic radius because the ionic force field diminishes with increasing radius, and consequently water molecules are less inclined to be abstracted from their position in bulk water.

Hydration numbers

Hydration numbers (the number of water molecules in the primary hydration layer) can be determined by various physical techniques (for example, compressibility) and the values obtained tend to differ depending on the method used. The overall total action of the ion on water may be replaced conceptually by a strong binding between the ion and some effective number (solvation number) of solvent molecules; this effective number may well be almost zero in the case of large ions

Hydrophobic hydration

Water is associated in a dynamic manner with nonpolar groups, but only in rare cases (where crystalline clathrates can be formed) is this water able to be isolated along with the hydrophobic groups. The phrase 'hydrophobic hydration' is used to describe this layer of water. The motion of water molecules is slowed down in the vicinity of nonpolar groups. Hydrophobic groups induce structure formation in water, hence the negative entropy (-A S) of their dissolution in water and the positive entropy (+AS) gained on their removal. In the discussion of hydrophobic bonding (Section 6.3.1) and nonpolar interactions, this special relationship between water and hydrocarbon chains is elaborated.

The solubility of inorganic materials in water

While a minority of therapeutic agents are inorganic electrolytes, it is nevertheless pertinent to consider the manner of their interaction with water. Electrolytes are, of necessity, components of replacement fluids, injections and eye drops and many other formulations. An increasing number of metal-containing compounds are used in diagnosis and therapy, some of which have interesting solution behaviour.

First consider the simpler salts. What determines the solubility of a salt such as sodium chloride and its solubility in relation to, say, silver chloride? The solubility of NaCl is in excess of 5 mol dm-3 while the solubility of AgCl is 500 000 times less. The heats of solution (AHsolution) are 62.8 kJ mol 1 for silver chloride and 4.2 kJ mol-1 for sodium chloride, suggesting a substantial difference either in the crystal properties or in the interaction of the ions with water. In fact the very great strength of the silver chloride crystal is due to the high polarisability of the silver ion. The heat of solution of an ionic solute can be written as

5.2.3 The effect of simple additives on solubility Solubility products

For poorly soluble materials such as silver chloride and barium sulfate the concept of the solubility product can be used. The following equilibrium exists in solution between crystalline silver chloride AgClc and ions in solution:

An equilibrium constant K can be defined as

[AgClJ V 7

Strictly, K should be written in terms of ther-modynamic activities and not concentrations, but activities can be replaced by concentrations (denoted by square brackets) because of the low solubilities involved (see section 3.3.1). At saturation the concentration of the crystalline silver chloride [AgClc] is essentially constant and the solubility product, Ksp, may therefore be written:

Conceptually, the solid salt (sodium chloride, for instance) is converted to the gaseous (g) state, Na+(g) + Cl(g), and each unit is then hydrated to form the species Na+(aq) and Cl (aq). If the heat of hydration is sufficient to provide the energy needed to overcome the lattice forces, the salt will be freely soluble at a given temperature and the ions will readily dislodge from the crystal lattice. If the partial molal enthalpy of solution of the substance is positive, the solubility will increase with increasing temperature; if it is negative the solubility will decrease, in agreement with Le Chatelier's principle.

The solubility product is useful for evaluating the influence of other species on the solubility of salts of low aqueous solubility. Some values of solubility products are quoted in Table 5.7.

Additives may either increase or decrease the solubility of a solute in a given solvent. The effect that they have will depend on several factors:

• The effect the additive has on the structure of water

• The interaction of the additive with the solute

• The interaction of the additive with the solvent

Table 5.7 Solubility products of some inorganic

salts

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