Atoms naturally seek maximum stability, which is achieved by having a full outer layer of electrons, known as the valence shell. This drive toward stability is the fundamental reason atoms interact, leading to the formation of chemical bonds. A chemical bond represents a powerful, lasting force of attraction that holds atoms together, transforming individual elements into compounds with new characteristics. These forces allow the diversity of matter to exist.
The Mechanism of Ionic Bonds
Ionic bonds form when there is a substantial difference in the atoms’ desire for electrons, typically occurring between a metal and a nonmetal. Metals readily give up their few valence electrons, while nonmetals easily accept electrons. This interaction involves a complete transfer of one or more electrons from the metal atom to the nonmetal atom.
The metal atom, upon losing an electron, becomes a positively charged ion called a cation. Conversely, the nonmetal atom gains the electron, becoming a negatively charged ion known as an anion. The bond itself is the strong electrostatic attraction between these oppositely charged ions. This attraction results in the ions packing together in a highly ordered, three-dimensional structure called a crystal lattice.
Common table salt, sodium chloride (\(\text{NaCl}\)), is the classic example. A sodium atom (\(\text{Na}\)) transfers its single valence electron to a chlorine atom (\(\text{Cl}\)). This forms a positive sodium ion (\(\text{Na}^+\)) and a negative chloride ion (\(\text{Cl}^-\)), which are locked into the rigid, repeating pattern of the salt crystal.
The Mechanism of Covalent Bonds
Covalent bonds typically form between two nonmetal atoms that have similar tendencies to attract electrons. Instead of surrendering electrons, the atoms achieve stability by sharing one or more pairs of valence electrons. The shared electrons are simultaneously attracted to the nuclei of both atoms, holding them together to form a discrete unit called a molecule.
The number of shared electron pairs determines the bond type: single (one pair), double (two pairs), or triple (three pairs). Unlike the extended crystal lattice of ionic compounds, covalent bonding results in distinct molecules, such as water (\(\text{H}_2\text{O}\)) or methane (\(\text{CH}_4\)). In a methane molecule, the carbon atom shares one electron pair with each of the four hydrogen atoms, forming a stable, tetrahedral molecule.
Resulting Properties of Bonded Compounds
The difference in bonding leads to contrasting physical properties. Ionic compounds, held together by the intense electrostatic force of the crystal lattice, exhibit high melting and boiling points. They exist as brittle solids at room temperature because any physical force that shifts the ions can align like-charged ions, causing repulsion and shattering the crystal.
Covalent compounds are composed of neutral molecules held together by much weaker intermolecular forces. Consequently, these compounds have lower melting and boiling points than their ionic counterparts, existing as solids, liquids, or gases at room temperature. Electrical conductivity is another differentiator: ionic compounds do not conduct electricity as solids but become excellent conductors when dissolved or melted, as their mobile ions are free to carry charge. Most covalent compounds do not conduct electricity in any state because they lack free-moving charged particles.
Understanding Polarity in Covalent Bonds
While covalent bonds involve electron sharing, this sharing is not always equal, introducing the concept of polarity. Unequal sharing is governed by electronegativity, an atom’s measure of its ability to attract the shared electrons in a bond toward itself. When two identical atoms bond, such as oxygen atoms in \(\text{O}_2\), the electronegativity difference is zero, resulting in a nonpolar covalent bond where electrons are shared equally.
When atoms of different elements bond, like hydrogen and oxygen in water, a difference in electronegativity exists. The electrons are pulled closer to the more electronegative atom (oxygen), giving it a partial negative charge (\(\delta^-\)). The less electronegative atom (hydrogen) acquires a partial positive charge (\(\delta^+\)). This creates a polar covalent bond, known as a bond dipole. Polarity dictates solubility; substances with similar polarity tend to mix, following the principle “like dissolves like.”