Ionic and Covalent Structures

Ionic Bonding and Structure

The force of attraction between oppositely charged ions (cations and anions) constitutes an ionic bond. Ionic bonding occurs between atoms of elements that have large differences in electronegativity, oftentimes between a metal with a very low electronegativity and a non-metal with a very high electronegativity (opp sides of the table). The units of ionic compounds such as sodium chloride and magnesium fluoride cannot be separated easily by direct heating of the crystal salts. The ions that make up the ionic solid are arranged in a specific array of repeating units which gives it a rigid structure, but the cations and anions are arranged so that the system has the minimum possible energy.

Magnesium fluoride has many properties that are characteristic of ionic solids. In general, ionic solids have the following properties:

• crystalline with smooth, shiny surfaces

• hard but brittle

• non-conductors of electricity and heat

• high melting points

Many ionic solids are also soluble in water. Magnesium fluoride, however, is an exception. It has a very low solubility in water.

Covalent Bonding and Structure

Covalent bonding involves a balance between the forces of attraction and repulsion that act between the nuclei and electrons of two or more atoms simultaneously. There is an optimum separation for two hydrogen atoms at which their nucleus-electron attractions, nucleus-nucleus repulsions, and electron-electron repulsions achieve this balance. This optimum separation favours a minimum energy for the system, and constitutes the covalent bond between the two hydrogen atoms. Unlike ionic bonding, in which electrons behave as if they are transferred from one atom to another, covalent bonding involves the sharing of pairs of electrons.

A covalent bond may form when two half-filled atomic orbitals from two atoms overlap to share the same region of space. A covalent bond involves the formation of a new orbital, caused by the overlapping of atomic orbitals. The new orbital has energy levels that are lower than those of the original atomic orbitals. Since electrons tend to occupy the lowest available energy level, the new orbitals provide a more energetically favourable configuration than the two atoms had before they interacted.

Bond energy is the energy required to break the force of attraction between two atoms in a bond and to separate them. Thus, bond energy is a measure of the strength of a bond. The bond energy increases if more electrons are shared between two atoms because there is an increase in charge density between the nuclei of the bonded atoms, so single bonds are weaker than double bonds.

In contrast to ionic solids, covalent (molecular) compounds typically have the following properties:

• exist as a soft solid, a liquid, or a gas at room temperature

• have low melting points and boiling points

• are poor conductors of electricity, even in solution

• may not be soluble in water (exceptions like acetone exist)

Metallic Bonding and Structure

About two-thirds of all the naturally occurring elements are metals. Despite this great diversity, metals have many properties in common. Metals conduct electricity and heat in both their solid and liquid states. Most metals are malleable and ductile; that is, they can be easily stretched, bent, and deformed without shattering of the whole solid, or can be easily made to be with the addition of heat or pressure. While metals show a broad range of melting and boiling points, in general, metals change state at moderate to high temperatures. Most metals have either one or two valence electrons.

Based on electronegativity differences, metals cannot form ionic bonds between themselves or other metals, and they also do not have enough valence electrons to form a covalent bond. Unlike the electron sharing in covalent compounds, however, electron sharing in metals occurs throughout the entire structure of the metal in a structure called the free-electron model, which pictures metals as being composed of a densely packed core of metallic cations, within a delocalized region of shared, mobile valence electrons.