Atoms seek stability by achieving a full outer shell of electrons, a concept known as the octet rule. The valence electrons involved in this process can be categorized based on their role in bonding. A standard covalent bond forms when two atoms share a pair of electrons, with each atom contributing one electron to that shared pair. Valence electrons that remain localized on a single atom are known as lone pairs or non-bonding pairs. This distinction raises the question of whether these unshared lone pairs can participate in forming a covalent bond.
The Non-Bonding Role of Lone Pairs
Lone pairs are valence electrons that are not shared with another atom and remain tightly held by the nucleus. These unshared pairs allow molecules to satisfy the octet rule without forming the maximum number of possible bonds. For example, the oxygen atom in water has two bonding pairs shared with hydrogen and two unshared lone pairs, fulfilling its octet.
Lone pairs significantly influence the three-dimensional shape of a molecule. According to Valence Shell Electron Pair Repulsion (VSEPR) theory, all electron pairs around a central atom repel each other, seeking maximum distance. Since a lone pair is attracted by only one nucleus, its electron cloud spreads out more broadly than a bonding pair, which is held by two nuclei.
This greater spatial distribution means a lone pair exerts a stronger repulsive force than a bonding pair. In the water molecule, the two lone pairs on the oxygen atom push the two hydrogen atoms closer together, bending the molecular shape from a predicted tetrahedral arrangement to a V-shape. Similarly, the single lone pair on nitrogen in ammonia (\(\text{NH}_3\)) pushes the three hydrogen atoms into a trigonal pyramidal shape. This repulsive effect is a primary function of lone pairs.
How Lone Pairs Form Coordinate Covalent Bonds
Lone pairs can form a specific type of covalent bond known as a coordinate covalent bond, or dative bond. The difference is that one atom supplies both electrons for the shared bond, rather than each atom supplying one. This process requires a donor atom, which possesses a lone pair, and an acceptor atom, which has a vacant orbital to accommodate the donated pair.
A classic example is the formation of the ammonium ion (\(\text{NH}_4^+\)) when ammonia (\(\text{NH}_3\)) reacts with a hydrogen ion (\(\text{H}^+\)). Nitrogen in ammonia has three covalent bonds and one lone pair. The hydrogen ion (\(\text{H}^+\)), which is a proton, has an empty valence orbital. Nitrogen acts as the donor, sharing its lone pair with the hydrogen ion acceptor, establishing a new coordinate covalent bond.
Another illustration is the formation of the hydronium ion (\(\text{H}_3\text{O}^+\)) when an acid dissolves in water. A water molecule has two lone pairs on its oxygen atom. One of these pairs can be donated to an electron-deficient hydrogen ion. This sharing of the oxygen lone pair with the empty orbital of the hydrogen ion creates the hydronium ion, a key species in aqueous acid solutions.
Covalent vs. Coordinate Bonds: A Comparison
The distinction between a standard covalent bond and a coordinate covalent bond lies solely in the origin of the shared electrons. In a typical covalent bond, both atoms contribute one electron to the shared pair. Conversely, in a coordinate covalent bond, one atom is the sole contributor, donating both electrons from its lone pair.
Despite this difference in formation, the resulting bonds are chemically indistinguishable once established. Once the coordinate bond is formed, it possesses the same properties as any other covalent bond in the molecule, including strength and length. For example, the four nitrogen-hydrogen bonds in the final ammonium ion are identical; it is impossible to tell which one was formed by the lone pair donation.
Both types of bonds involve the sharing of two electrons between two atoms. The terminology merely serves to describe the specific mechanism by which the bond initially formed.