Chemical bonds are the attractive forces that hold atoms together to form compounds and molecules. Atoms form these bonds to achieve a more stable, lower-energy state, usually by completing their outermost electron shells. The two primary categories of chemical bonding are ionic and covalent, distinguished by how participating atoms handle their valence electrons.
How Ionic Bonds Form
Ionic bonds form through the complete transfer of one or more valence electrons from one atom to another. This transfer occurs between a metal atom, which readily loses electrons, and a nonmetal atom, which readily gains electrons. The atom that loses electrons becomes a positively charged ion (a cation), and the atom that gains electrons becomes a negatively charged ion (an anion).
A classic example is the formation of sodium chloride (\(\text{NaCl}\)), or table salt, where the sodium atom transfers its single valence electron to the chlorine atom. Sodium becomes a stable \(\text{Na}^+\) cation, and chlorine becomes a stable \(\text{Cl}^-\) anion, both achieving a full outer electron shell. The final bond is not a shared pair of electrons, but the powerful electrostatic attraction between these oppositely charged ions. This attraction pulls the ions together into a repeating, three-dimensional structure known as a crystal lattice.
How Covalent Bonds Form
Covalent bonds result from the mutual sharing of valence electrons between two atoms, most commonly occurring between two nonmetal atoms. Instead of a transfer, the atoms pool their electrons into a shared region of space between their nuclei. This sharing allows both atoms to count the shared electrons toward completing their outer electron shells.
The bond itself is the attractive force exerted by both atoms’ nuclei on the shared electron pair located in the space between them. For instance, in a water molecule (\(\text{H}_2\text{O}\)), the central oxygen atom shares electrons with two hydrogen atoms. In the oxygen molecule (\(\text{O}_2\)), two oxygen atoms share two pairs of electrons, forming a double bond.
The Electronegativity Continuum
The fundamental difference between these two bond types is best understood as a spectrum, or continuum, governed by a property called electronegativity. Electronegativity is defined as an atom’s inherent ability to attract a shared pair of electrons toward itself within a chemical bond. The difference in electronegativity (\(\Delta \text{EN}\)) between the two bonded atoms determines the nature of the bond.
A very large difference in electronegativity signifies that one atom has a significantly stronger pull on the electrons, leading to the complete electron transfer characteristic of an ionic bond. Numerical guidelines suggest that a \(\Delta \text{EN}\) greater than 1.7 or 2.0 on the Pauling scale indicates the bond is predominantly ionic. Conversely, when the electronegativity difference is small or zero, the electrons are shared relatively equally, resulting in a nonpolar covalent bond, such as in \(\text{O}_2\).
Bonds with an intermediate \(\Delta \text{EN}\) (between 0.5 and 1.7) are classified as polar covalent, where electrons are shared unequally. In these bonds, the electron pair spends more time closer to the more electronegative atom, creating partial negative and partial positive charges. The transfer is not complete enough, however, to be considered fully ionic.
Contrasting Physical Characteristics
The distinct mechanisms of ionic transfer versus covalent sharing lead directly to differences in the resulting compounds’ physical properties. Ionic compounds are held together by strong, non-directional electrostatic forces within a lattice structure, requiring substantial energy to break apart. Consequently, they exhibit very high melting and boiling points, and are hard, brittle solids at room temperature.
Covalent compounds, in contrast, are formed by discrete molecules held together by much weaker intermolecular forces. Breaking these weak forces to cause a phase change requires far less energy, resulting in significantly lower melting and boiling points than ionic compounds. For this reason, covalent substances can exist as gases (like \(\text{O}_2\)), liquids (like \(\text{H}_2\text{O}\)), or soft solids at standard conditions.
Furthermore, the electrical conductivity of the two compound types differs greatly. Ionic compounds do not conduct electricity in their solid state because the charged ions are locked in the rigid crystal lattice. When dissolved in water or melted, the ions become mobile, allowing the compound to conduct an electrical current. Covalent compounds, which are composed of neutral molecules rather than mobile ions, do not conduct electricity in any state.