Ionic compounds are predominantly formed by the chemical combination of a metal and a nonmetal. This relationship is based on the opposing electronic tendencies of these two classes of elements. Ionic compounds are defined by the electrostatic attraction between oppositely charged ions, which results from a complete transfer of electrons from a metallic atom to a nonmetallic atom. Common table salt, or sodium chloride (\(\text{NaCl}\)), serves as the classic example of this interaction.
The Elemental Difference: Metals and Nonmetals
The propensity for ionic bond formation begins with the distinct positions of metals and nonmetals on the Periodic Table. Metals are found on the left side and center and possess low ionization energy. This means little energy is required for a metallic atom to lose its outermost valence electrons. These atoms, such as those in Groups 1 and 2, readily surrender their electrons to achieve a stable outer shell, forming positively charged ions known as cations.
Nonmetals, located on the right side of the table (excluding the noble gases), exhibit the opposite electronic behavior. They have high electronegativity, which describes a strong attraction for electrons in a chemical bond. Nonmetallic atoms, particularly those in Groups 16 and 17, seek to gain electrons to fill their own valence shells. By accepting electrons, they become negatively charged ions called anions.
The large difference in electronegativity between a metal and a nonmetal is the driving force behind ionic bonding. The metal atom has a weak hold on its valence electrons, while the nonmetal atom has a powerful pull for them. This contrast ensures that electrons are completely transferred, rather than being shared, forming the basis of the electrostatic attraction.
The Mechanism of Electron Transfer
The formation of the ionic bond is driven by the atoms’ desire to achieve the stability of a noble gas, a concept often simplified as the Octet Rule. Noble gases have a full complement of eight electrons in their outermost shell, a highly stable arrangement. To mimic this, a metal like sodium, which has one valence electron, easily gives it up.
When sodium loses this single electron, it becomes the positively charged sodium ion (\(\text{Na}^+\)). Simultaneously, a nonmetal like chlorine, which has seven valence electrons, readily accepts that electron. By gaining one electron, chlorine achieves a full octet and becomes the negatively charged chloride ion (\(\text{Cl}^-\)).
This complete transfer of electrons differentiates ionic bonding from covalent bonding, where electrons are shared. The resulting ions, \(\text{Na}^+\) and \(\text{Cl}^-\), are powerfully drawn to each other by strong electrostatic forces. This attraction holds the ions together in a stable, three-dimensional arrangement called a crystal lattice. The energy released during the formation of this structure is referred to as lattice energy.
Defining Characteristics of Ionic Compounds
The strong electrostatic forces within the crystal lattice impart specific, observable physical characteristics to ionic compounds. One prominent feature is their high melting and boiling points. Significant thermal energy is required to overcome the strong attraction holding the ions together in the lattice structure. For instance, sodium chloride melts at about \(801^\circ\text{C}\).
Ionic compounds are hard and brittle as solids. If a mechanical force is applied to the crystal, it causes one layer of ions to shift relative to another. This movement brings ions with the same charge into direct alignment, leading to a repulsive force that shatters the crystal structure along a plane.
A defining property is their electrical conductivity. In their solid state, ionic compounds do not conduct electricity because the ions are locked rigidly in place within the lattice. However, when dissolved in a polar solvent like water or heated until they melt, the ions are released. The mobile, charged ions are then free to move, allowing the solution or molten compound to conduct an electric current.
Ionic compounds tend to be soluble in polar solvents, such as water. The polar water molecules are effective at surrounding and separating the oppositely charged ions in the lattice. This allows the compound to dissociate into its constituent ions.