- Related Topics:
- liquid crystal
- solution
- pure liquid
- coalescence
- liquid mixture
- On the Web:
- PNAS - Significant structures in the liquid state (PDF) (Mar. 29, 2025)
For those solutions in which there are strong intermolecular forces due to large dipole moments, hydrogen bonding, or complex formation, equations based on fundamental molecular theory cannot be applied, but it is frequently useful to apply a chemical treatment—i.e., to describe the liquid mixture in terms of association and solvation, by assuming the existence of a variety of distinct chemical species in chemical equilibrium with one another. For example, there is much experimental evidence for association in acetic acid, in which most of the molecules dimerize; i.e., two single acetic acid molecules, called monomers, combine to form a new molecule, called a dimer, through hydrogen bonding. When acetic acid is dissolved in a solvent such as benzene, the extent of dimerization of acetic acid depends on the temperature and on the total concentration of acetic acid in the solution. The escaping tendency (vapour pressure) of a monomer is much greater than that of a dimer, and it is thus possible to explain the variation of activity coefficient with composition for acetic acid in benzene; the activity coefficient of acetic acid in an excess of benzene is large because, under these conditions, acetic acid is primarily in the monomeric state, whereas pure acetic acid is almost completely dimerized. In the acetic acid–benzene system, association of acetic acid molecules produces positive deviations from Raoult’s law.
When a solvent and a solute molecule link together with weak bonds, the process is called solvation. For example, in the system acetone-chloroform, a hydrogen bond is formed between the hydrogen atom in chloroform and the oxygen atom in acetone. In this case, hydrogen bonding depresses the escaping tendencies of both components, producing negative deviations from Raoult’s law.
While hydrogen bonding is frequently encountered in solutions, there are many other examples of weak chemical-bond formation between dissimilar molecules. The formation of such weak bonds is called complex formation—that is, formation of a new chemical species, called a complex, which is held together by weak forces that are chemical in nature rather than physical. Such complexes usually exist only in solution; because of their low stability, they cannot, in general, be isolated. The ability of molecules to form complexes has a strong effect on solution behaviour. For example, the solubility of a sparingly soluble species can be much increased by complex formation: the solubility of silver chloride in water is extremely small since silver chloride dissociates only slightly to silver ion and chloride ion; however, when a small quantity of ammonia is added, solubility rises dramatically because of the reaction of six molecules of ammonia with one silver ion to form the complex ion Ag(NH3)6+. By tying up silver ions and forcing extensive dissociation of molecular silver chloride, the ammonia pulls silver chloride into aqueous solution.
In recent years there has been much interest in the use of computers to generate theoretical expressions for the activity coefficients of solutions. In many cases the calculations have been restricted to model systems, in particular to mixtures of hard-sphere (envisioned as billiard balls) molecules—i.e., idealized molecules that have finite size but no forces of attraction. These calculations have produced a better understanding of the structure of simple liquid solutions since the manner in which nonpolar and non-hydrogen-bonding molecules arrange themselves in space is determined primarily by their size and shape and only secondarily by their attractive intermolecular forces. The results obtained for hard-sphere molecules can be extended to real molecules by applying corrections required for attractive forces and for the “softness” of the molecules—i.e., the ability of molecules to interpenetrate (overlap) at high temperatures. While practical results are still severely limited and while the amount of required computer calculation is large even for simple binary systems, there is good reason to believe that advances in the theory of solution will increasingly depend on computerized, as opposed to analytical, models.
Solutions of electrolytes
Near the end of the 19th century, the properties of electrolyte solutions were investigated extensively by the early workers in physical chemistry. A suggestion of Svante August Arrhenius, a Swedish chemist, that salts of strong acids and bases (for example, sodium chloride) are completely dissociated into ions when in aqueous solution received strong support from electrical-conductivity measurements and from molecular-weight studies (freezing-point depression, boiling-point elevation, and osmotic pressure). These studies showed that the number of solute particles was larger than it would be if no dissociation occurred. For example, a 0.001 molal solution of a uni-univalent electrolyte (one in which each ion has a valence, or charge, of 1, and, when dissociated, two ions are produced) such as sodium chloride, Na+Cl-, exhibits colligative properties corresponding to a nonelectrolyte solution whose molality is 0.002; the colligative properties of a 0.001 molal solution of a univalent-divalent electrolyte (yielding three ions) such as magnesium bromide, Mg2+Br2-, correspond to those of a nonelectrolyte solution with a molality of 0.003. At somewhat higher concentrations the experimental data showed some inconsistencies with Arrhenius’ dissociation theory, and initially these were ascribed to incomplete, or partial, dissociation. In the years 1920–30, however, it was shown that these inconsistencies could be explained by electrostatic interactions (Coulomb forces) of the ions in solution. The current view of electrolyte solutions is that, in water at normal temperatures, the salts of strong acids and strong bases are completely dissociated into ions at all concentrations up to the solubility limit. At high concentrations Coulombic interactions may cause the formation of ion pairs, which implies that the ions are not dispersed uniformly in the solution but have a tendency to form two-ion aggregates in which a positive ion seeks the close proximity of a negative ion and vice versa. While the theory of dilute electrolyte solutions is well advanced, no adequate theory exists for concentrated electrolyte solutions primarily because of the long-range Coulomb forces that dominate in ionic solutions.
The equilibrium properties of electrolyte solutions can be studied experimentally by electrochemical measurements, freezing-point depressions, solubility determinations, osmotic pressures, or measurements of vapour pressure. Most electrolytes, such as salts, are nonvolatile at ordinary temperature, and, in that event, the vapour pressure exerted by the solution is the same as the partial pressure of the solvent. The activity coefficient of the solvent can, therefore, be found from total-pressure measurements, and, using the Gibbs-Duhem equation, it is then possible to calculate the activity coefficient of the electrolyte solute. This activity coefficient is designated by γ± to indicate that it is a mean activity coefficient for the positive and negative ions. Since it is impossible to isolate positive ions and negative ions into separate containers, it is not possible to determine individual activity coefficients for the positive ions and for the negative ions. The mean activity coefficient γ± is so defined that it approaches a value of unity at very low molality where the ions are so far apart that they exert negligible influence on one another. For small concentrations of electrolyte, the theory of Peter Debye, a Dutch-born American physical chemist, and Erich Hückel, a German chemist, relates γ± to the ionic strength, which is the sum of the products of the concentration of each ion (in moles per litre) and the square of its charge; the equation predicts that γ± decreases with rising ionic strength in agreement with experiment at very low ionic strength; at higher ionic strength, however, γ± rises, and in some cases γ± is greater than 1. The derivation of the Debye-Hückel theory clearly shows that it is limited to low concentrations. Many attempts have been made to extend the Debye-Hückel equation to higher electrolyte concentrations. One of the more successful attempts is based on the idea that the ions are solvated, which means that every ion is surrounded by a tight-fitting shell of solvent molecules.
The concept of solvation is often used to explain properties of aqueous solutions; one well-known property is the salting-out effect, in which the solubility of a nonelectrolyte in water is decreased when electrolyte is added. For example, the solubility of ethyl ether in water at 25° C is 0.91 mole percent, but, in an aqueous solution containing 15 weight percent sodium chloride, it is only 0.13 mole percent. This decrease in solubility can be explained by postulating that some of the water molecules cannot participate in the dissolution of the ether because they are tightly held (solvated) by sodium and chloride ions.
Electrolyte solutions have long been of interest in industry since many common inorganic chemicals are directly obtained, or else separated, by crystallization from aqueous solution. Further, many important chemical and metallurgical products (e.g., aluminum) are obtained or refined by electrochemical processes that occur in liquid solution. In recent years there has been renewed interest in electrolyte solutions because of their relevance to fuel cells as a possible source of power for automobiles.
The properties of electrolyte solutions also have large importance in physiology. Many molecules that occur in biological systems bear electric charges; a large molecule that has a positive electric charge at one end and a negative charge at the other is called a zwitterion. Very large molecules, such as those of proteins, may have numerous positive and negative charges; such molecules are called polyelectrolytes. In solution, the conformation (i.e., the three-dimensional structure) of a large, charged molecule is strongly dependent on the ionic strength of the dissolving medium; for example, depending on the nature and concentration of salts present in the solvent, a polyelectrolyte molecule may coagulate into a ball, it may stretch out like a rod, or it may form a coil or helix. The conformation, in turn, is closely related to the molecule’s physiological function. As a result, improved understanding of the properties of electrolyte solutions has direct consequences in molecular biology and medicine.
