Dissociation

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Dissociation

Dissociation is the process by which a chemical combination breaks up into simpler constituents. Dissociation may be accomplished by the addition of energy, as in the case of gaseous molecules dissociated by heat; or by the action of a solvent on a polar compound (i.e., electrolytic decomposition). All electrolytes dissociate to some extent in polar solvents. The degree of dissociation can be used to calculate the equilibrium constant for the dissociation reaction.

In the case of diatomic molecules, the dissociation energy refers to the energy required to break the gaseous molecules into their constituent atoms. Generally, for families of molecules such as HF, HCl, HBr, and HI; or C₂, Si₂, and Pb₂, the heavier molecules tend to have smaller dissociation energies. In the case of N₂, O₂, and F₂, the dissociation energy increases with increasing net bonding. In general, there is an inverse relationship between dissociation energy and bond length for related molecules. In the case of polyatomic molecules, resulting in the formation of molecular fragments, the energy required to break a single bond is referred to as a bond dissociation energy.

When the acid hydrogen chloride (HCl) dissolves in water, a dipole-dipole interaction occurs between the negative dipole of one molecule and the positive side the other. The water molecules pull on the HCl molecules, tending to dissociate them into H⁺ ions and Cl^- ions. Opposing this tendency is the covalent bond holding the HCl molecule together. In the case of the water-HCl interaction, the dissociation reaction predominates, and H⁺ ions are formed in the solution. The production of H⁺ ions in aqueous solutions is a property common to all acids. Bases, on the other hand, produce OH^- ions in water, forming basic (or alkaline) solutions.

Acids and bases, for all practical purposes, exist, because water dissociates into H⁺ and OH^- ions. In equilibrium, only a few water molecules are dissociated. The general law of equilibrium predicts an equilibrium constant for this dissociation of [H⁺][OH^-]/H₂O], but since the concentration of water in aqueous solutions is very large, this constant is usually put into the form [H⁺][OH^-] = 1.0 x 10^-14 at 25 deg C. In plain language, this expression says that any water solution at 25 deg C will contain H⁺ and OH^- ions in such concentrations that their product will be 1.0 x 10^-14.

Ordinarily the concentrations of H⁺ and OH^- ions in a solution are not equal. A water solution in which the H⁺ ion concentration exceed the OH^- concentration is called acidic; solutions in which the OH^- concentration is larger are basic.

Weak acids and bases form H⁺ and OH^- ions, respectively, in solution only to a small extent. In the case of the hydrogen fluoride (HF), the acid dissociation constant is defined as [H⁺][F^-]/HF]. For ammonia (NH₃; a base) in water, the base dissociation constant is [H⁺][OH^-]/NH₃]. When a compound containing a particular complex ion is added to water, it is customary to treat the ensuing reaction as if it were a simple dissociation. The dissociation constant of the complex ion is written analogously.