Solvation of Ions

Reactions which involve the formation of charged atoms and molecules are usually extremely endothermic in the gas phase, but may become spontaneous in certain solvents. If ions are formed from a neutral compound, as when NaCl is dissolved in water, the oppositely charged cations and anions naturally attract each other, so formation of a dispersed homogeneous solution might appear to be energetically unfavorable. To achieve charge separation of ions in solution, two solvent characteristics are particularly important. The first is the ability of solvent molecules to orient themselves between ions so as to attenuate the electrostatic force one ion exerts on the other. This characteristic is a function of the polarity of the solvent. Solvent polarity has been defined and measured in several different ways, one of the most common being the dielectric constant, ε. High dielectric constant solvents such as water (ε=80), dimethyl sulfoxide (ε=48) & N,N-dimethylformamide (ε=39), usually have polar functional groups, and often high dipole moments. When subject to the electric field of an ion, such polar molecules orient themselves to oppose the field, and in so doing they limit its reach. Because of electrostatic attraction between these polar groups, the boiling points of these solvents are generally higher than those of similarly sized nonpolar solvents, such as diethyl ether (ε=4.3) and hexane (ε=1.9).
Solvents that have relatively acidic hydrogen atoms (e.g. O-H & N-H) are called protic. Because their functional groups are made up of polar covalent bonds, protic solvents are often polar as well. A list of common protic and aprotic solvents is provided here. The dielectric constants provide a measure of solvent polarity.

The second factor important in the stabilization of ions, which also resists their intimate recombination, is called solvation. This refers to the ability of solvent molecules to stabilize ions by encasing them in a sheath of weakly bonded solvent molecules, thus somewhat dispersing the electrical charge. Anions are best solvated by hydrogen-bonding solvents; cations are generally solvated by binding to nucleophilic sites on a solvent molecule Two-dimensional diagrams illustrating the primary solvation shell about Na⁽⁺⁾ and Cl^(–) are shown here. The water dipoles are drawn as red arrows, and partial charges are noted. Additional water molecules are oriented in secondary and tertiary layers about the ions.

From this description of ion formation in solution, it should be clear that both enthalpy and entropy factors will be important to the outcome of an ionization process. Thus solvation stabilizes and insulates an ion, helping the enthalpic change, whereas the same solvation adds order and structure to the ionic species at the cost of lowering entropy. The outcome of these interactions is discussed below for two typical salts.

Although these two inorganic salts have similar standard enthalpies of solution in water, their standard entropies are quite different. One might expect this entropy change to be positive, since a single molecule in the solid state produces two or more ionic species, accompanied by an increase in system disorder. However this argument fails to consider the ordering of solvent molecules taking place in the solvation of these ions. Because of their greater charge density, small ions and highly charged ions, such as F^– and Ca²⁺, require greater solvation than large or singly charged ions, such as Na⁺ or Cl^–. The overall entropy change for solution of NaCl is positive, reflecting the increased disorder of ionization, but the entropy change for CaF₂ solution is strongly negative thanks to the solvation shell structure required by the resulting ions. These different entropy changes are incorporated in the free energy of solution, which is exergonic for NaCl, but endergonic for CaF₂. The result is dramatic. Sodium chloride is quite soluble in water at room temperature (36g per 100g water), but calcium fluoride is nearly insoluble (0.0016g per 100g water).