Ionic Bonds

Although atoms with equal numbers of protons and electrons exhibit no electrical charge, it is common for atoms to attain the stable electronic configuration of the inert gases by either gaining or losing electrons. The metallic elements on the left side of the periodic table have electrons in excess of the stable configuration. Table 1 shows the electron loss necessary for three light metals to reach a stable electron structure.

The positive charge on the resulting metal ion is due to the atom possessing more nuclear protons than orbital electrons. The valence electrons are most distant from the nucleus; thus, they are weakly held by the electrostatic attraction of the protons and, consequently, are easily stripped from atoms of the metals.

By contrast, the nonmetallic elements on the right side of the periodic table have fewer electrons than that of a stable configuration and can most readily attain the stable configuration of the inert gases by gaining electrons. The negative charge on the resulting nonmetal ion is due to the atom possessing more orbital electrons than nuclear protons. Table 2 compares three nonmetals to the inert gas argon.

Because metallic elements tend to lose electrons and nonmetallic elements tend to gain electrons, a pair of contrasting elements will exchange electrons so that both achieve stable electronic configurations. The resulting ions of opposite charge have a strong force of electrostatic attraction, which is called an ionic bond. Note: This bond forms through the complete transfer of electrons from one atom to another, in contrast to the electron sharing of the covalent bond.

The force of attraction between two points of opposite electrical charge is given by Coulomb's law: 

where q + is the positive charge, q is the negative charge, and d is the distance between the two charges. This law of electrostatic attraction can be used to measure the distance between two spherical ions because the charges can be considered to be located at the center of each sphere. (See Figure 1.)

Figure 1. The distance between ionic charges.


Notice that the distance between the centers of the two ions is the sum of radii of the ions. The appropriate electrostatic force then is calculated from the equation 

where q C is the charge of the positive cation, q A is the charge of the negative anion, and the denominator is the sum of their radii.

The strength of ionic bonding, therefore, depends on both the charges and the sizes of the two ions. Higher charges and smaller sizes produce stronger bonds. Table 3 shows the approximate radii of selected ions, which have the electronic configuration of an inert gas. The radii are in .

For ions of the same charge, the ionic radius increases as you go down any column because the elements of higher atomic number have a greater number of electrons in a series of electronic shells progressively farther from the nucleus. The change in ionic size along a row in the chart just above shows the effect of attraction by protons in the nucleus.

In Table 4, the five ions O 2– through A1 3+ are all isoelectronic; that is, they have the same number of electrons in the same orbitals.

For continuity, the neutral Ne atom is also in the chart, with its atomic radius. As you proceed to the right in Table 4, the greater number of protons attracts the electrons more strongly, producing progressively smaller ions.

Now use Coulomb's law to compare the strengths of the ionic bonds in crystals of magnesium oxide and lithium fluoride. The sizes of the four ions are taken from the tabulation of radii of cations and anions in Table 5.

Comparing the two relative forces of electrostatic attraction that you calculated, you can conclude that ionic bonding is considerably stronger in magnesium oxide. This affects the physical properties and chemical behavior of the two compounds. For example, the melting point of MgO (2,852°C) is much higher than that of LiF (845°C).

The strength of chemical bonding in various substances is commonly measured by the thermal energy (heat) needed to separate the bonded atoms or ions into individual atoms or ions.