Mechanism of solute-solvent interaction
A solute dissolves in a solvent when it forms favourable interactions with the solvent. This dissolving process all depends upon the free energy changes of both solute and solvent. The free energy of solvation is a combination of several factors. The process can be considered in three stages:
(i) A solute (drug) molecule is ‘removed’ from its crystal.
The solute must separate out from the bulk solute. This is enthalpically unfavourable as solute-solute interactions are breaking but are entropically favourable.
(ii) A cavity for the drug molecule is created in the solvent.
A cavity must be created in the solvent. The creation of the cavity will be entropically and enthalpically unfavourable as the ordered structure of the solvent decreases and there are fewer solvent-solvent interactions.
(iii) The solute (drug) molecule is inserted into this cavity.
The solute must occupy the cavity created in the solvent. Placing the solute molecule in the solvent cavity requires a number of solute-solvent contacts; the larger the solute molecule, the more contacts are created. If the surface area of the solute molecule is A, and the solute-solvent interface increases by γ12A, where γ12is the interfacial tension between the solvent 1and the solute 2 then it leads to favourable solute-solvent interactions. This is entropically favourable as the mixture is more disordered than when the solute and solvent are not mixed.
Dissolution often occurs when the solute-solvent interactions are similar to the solvent-solvent interactions, signified by the term ‘Like dissolves Like’. Hence, polar solutes dissolve in polar solvents, whereas non-polar solutes dissolve in non-polar solvents. The dissimilar nature of solute and solvent makes solute insoluble in the solvent. Substances dissolve when the solvent-solute attraction is greater than solvent-solvent attraction and solute-solute attraction.
- The hole or cavity of in the solvent created in step (b) is now closed, and an additional decrease in energy, —w )2 occurs, involving net work in this final step of —w12. The total work as given by this extremely simplified scheme is thus
w22 + w11 — 2W12
- The activity coefficient term of the solubility equations however, was shown by Scatchard and by Hildebrand and Wood19 to be proportional also to the volume of the solute, considered as a supercooled liquid, and to the fraction of the total volume occupied by the solvent. The logarithm of the activity coefficient is given by the more elaborate expression
- where V2 is the molar volume or volume per mole of (supercooled) liquid solute and (φ 1 is the volume fraction, or X1 V1/(X1V1 + X2V2), of the solvent. R is the gas constant 1.987 cal/mole deg, and T is the absolute temperature of the solution.
The pharmacist knows that water is a good solvent for salts, sugars, and similar compounds, whereas mineral oil and benzene are often solvents for substances that are normally only slightly soluble in water. These empirical findings are summarized in the statement, “like dissolves like.” Such a maxim is satisfying to most of us, but the inquisitive student may be troubled by this vague idea of “likeness.” The student who sets out to learn in what manner the solute and solvent are alike will discover a fascinating area of scientific investigation that is still in an unsettled state. The advanced student interested in this subject should consult the work of Hildebrand and Scott,3 Leussing,4 and Dack,5
The solubility of a drug is due in large measure to the polarity of the solvent, that is, to its dipole moment. Polar solvents dissolve ionic solutes and other polar substances. Accordingly, water mixes in all proportions with alcohol and dissolves sugars and other polyhydroxy compounds.
Hildebrand showed, however, that consideration of dipole moments alone is not adequate to explain the solubility of polar substances in water. The ability of the solute to form hydrogen bonds is a far more significant factor than is the polarity as reflected in a high dipole moment. Although nitrobenzene has a dipole moment of 4.2 x 10-18 esu cm, and phenol has a value of only 1.7 x 10—18 esu cm, nitrobenzene is soluble only to the extent of 0.0155 mole/kg in water. Whereas phenol is soluble to the extent of 0.95 mole/kg at 20°C.
Water dissolves phenols alcohols, aldehydes, ketones, alflines, and other oxygen- and nitrogen-containing compounds that can form hydrogen bonds with water:
A difference in the acidic and basic character of the constituents in the Lewis electron donor-acceptor sense also contributes to specific interactions in solutions.
The molecules of water in ice are joined together by hydrogen bonds to yield a tetrahedral structure. Although some of the hydrogen bonds are broken when the ice melts, water retains its ice-like structure in large measure at ordinary temperatures. This quasicrystalline structure is broken down when water is mixed with another substance that is capable of hydrogen bonding. When ethyl alcohol and water are mixed, the hydrogen bonds between the water molecules are replaced in part by hydrogen bonds between water and alcohol molecules.
In addition to the factors already enumerated, the solubility of a substance also depends on structural features such as the ratio of the polar to the nonpolar groups of the molecule. As the length of a nonpolar chain of an aliphatic alcohol increases, the solubility of the compound in water decreases. Straight-chain monohydroxy alcohols, aldehydes, ketones, and acids with more than four or five carbons cannot enter into the hydrogen-bonded structure of water and hence are only slightly soluble. When additional polar groups are present in the molecule, as found in propylene glycol, glycerin, and tartaric acid, water solubility increases greatly. Branching of the carbon chain reduces the nonpolar effect and leads to increased water solubility. Tertiary butyl alcohol is miscible in all proportions with water, whereas n-butyl alcohol dissolves to the extent of about 8 g/100 mL of water at 20°C.
In brief, polar solvents such as water act as solvents according to the following mechanisms6:
(a) Owing to their high dielectric constant, namely about 80 for water, polar solvents reduce the force of attraction between oppositely charged ions in crystals such as sodium chloride (p. 34). Chloroform has a dielectric constant of 5 and benzene that of about 2; hence, ionic compounds are practically insoluble in these solvents.
(b) Polar solvents break covalent bonds of potentially strong electrolytes by acid-base reactions because these solvents are amphiprotic (p. 162). For example, water brings about the ionization of HC1 as follows:
HCl + H20 H3O+ + Cl–
Weak organic acids are not ionized appreciably by water; their partial solubility is attributed instead to the hydrogen bond formation with water. Phenols and carboxylic acids, however, are readily dissolved in solutions of strong bases.
(c) Finally, polar solvents are capable of solvating molecules and ions through dipole interaction forces, particularly hydrogen bond formation, which leads to the solubility of the compound. The solute must be polar in nature because it often must compete for the bonds of the already associated solvent molecules if it is to win a place in the associated structure. The ion-dipole interaction between the sodium salt of oleic acid and water can be depicted as
The solvent action of nonpolar liquids, such as hydrocarbons, differs from that of polar substances. Nonpolar solvents are unable to reduce the attraction between the ions of strong and weak electrolytes because of the solvents’ low dielectric constants. Nor can the solvents break covalent bonds and ionize weak electrolytes, because they belong to the group known as aprotic solvents (p. 161), and they cannot form hydrogen bridges with nonelectrolytes. Hence, ionic and polar solutes are not soluble or are only slightly soluble in nonpolar solvents.
Nonpolar compounds, however, can dissolve nonpolar solutes with similar internal pressures (p. 244) through induced dipole interactions. The solute molecules are kept in solution by the weak van der Waals—London type of forces (p. 23) Thus, oils and fats dissolve in carbon tetrachloride, benzene and mineral oil. Alkaloidal bases and fatty acids also dissolve in nonpolar solvents.
Semipolar solvents, such as ketones and alcohols, can induce a certain degree of polarity in nonpolar solvent molecules, so that, for example, benzene, which is readily polarizable, becomes soluble in alcohol. In fact, semipolar compounds can act as intermediate solvents to bring about the miscibility of polar and nonpolar liquids. Accordingly, acetone increases the solubility of ether in water. Loran and Guth studied the intermediate solvent action of alcohol on the water—castor oil mixtures. Propylene glycol has been shown to increase the mutual solubility of water and peppermint oil and of water and benzyl benzoate.
A number of common solvent types are listed in the order of decreasing “polarity” in Table 10—2, together with corresponding solute classes. The term polarity is loosely used here to represent not only the dielectric constants of the solvents and solutes but also the other factors enumerated previously.
Mr. Sumeet Manohar Kharat
PSPS’s Indira Institute of Pharmacy, Sadavali
Physical Pharmaceutics I
SOLUTE SOLVENT INTERACTIONS VIDEO LECTURE
Video Lecture Credit:
Dr. Govind Kailash Lohiya
Gurunanak College of Pharmacy, Nagpur
Physical Pharmaceutics I
Solubility of drugs Physical Pharmaceutics Mechanism of solute solvent interaction