The octet rule is a foundational concept in chemistry that explains the bonding behavior of many atoms. This rule posits that atoms tend to gain, lose, or share electrons until their outer shell contains eight valence electrons, achieving a stable configuration similar to that of noble gases. While this generalization provides a simple framework for understanding how molecules form, it is not a universal law. Certain elements and molecules deviate from this pattern, creating exceptions that fall into three main categories: incomplete octets, expanded octets, and molecules with an odd total number of electrons.
Molecules with Incomplete Octets
Molecules with incomplete octets are those where the central atom is stable with fewer than eight valence electrons. This stability is most commonly observed in compounds involving Group 13 elements, specifically Boron (B) and Aluminum (Al), as well as Beryllium (Be). These atoms naturally have few valence electrons, making it favorable to bond with fewer partners rather than striving for a full octet.
Boron, for instance, has three valence electrons. In molecules like Boron trifluoride (\(\text{BF}_3\)), it forms three single bonds, leaving the central boron atom with six valence electrons. Attempting to achieve an octet by forming a double bond would result in highly unfavorable formal charges. Therefore, the molecule remains stable in its electron-deficient state, preferring the incomplete octet to the energetic cost of a complete one.
These electron-deficient species are highly reactive because they possess empty orbitals eager to accept an electron pair. This characteristic classifies \(\text{BF}_3\) as a powerful Lewis acid, meaning it readily accepts electrons from another molecule to complete its octet. Beryllium, with two valence electrons, is another example; in Beryllium dichloride (\(\text{BeCl}_2\)), the central atom achieves stability with only four valence electrons.
Molecules with Expanded Octets
Expanded octets occur when the central atom accommodates more than eight valence electrons. This phenomenon is only possible for nonmetal elements located in the third period and below, such as Phosphorus (P), Sulfur (S), and Iodine (I). Second-period elements like Nitrogen and Oxygen cannot exhibit this behavior because they only have the 2s and 2p orbitals available for bonding, which hold a maximum of eight electrons.
Atoms in the third period and subsequent periods have accessible d-orbitals in their valence shell that can participate in bonding. Although these d-orbitals are typically unoccupied in the ground state, they provide the “extra space” needed to house additional electron pairs. For example, in Phosphorus pentachloride (\(\text{PCl}_5\)), the central phosphorus atom is surrounded by five chlorine atoms, bonding it to ten valence electrons.
Sulfur hexafluoride (\(\text{SF}_6\)) provides an example where the central sulfur atom forms six bonds, resulting in twelve valence electrons. The ability to expand the octet is tied to minimizing formal charges within the structure. Drawing a Lewis structure that strictly adheres to the octet rule often results in significant, energetically unfavorable charge separation. The molecule instead utilizes available d-orbitals to form additional bonds, distributing electrons and bringing formal charges closer to zero, resulting in a more stable configuration.
Odd-Electron Molecules
The third exception involves molecules that possess an odd total number of valence electrons. Since electrons exist in pairs in covalent bonds and lone pairs, it is mathematically impossible to distribute an odd number of electrons so that every atom achieves a full octet. One atom must be left with an unpaired electron, meaning it has only seven electrons surrounding it.
These molecules are commonly referred to as free radicals. Examples include Nitric Oxide (\(\text{NO}\)), which has eleven valence electrons, and Nitrogen Dioxide (\(\text{NO}_2\)), which has seventeen valence electrons. The unpaired electron is often placed on the atom that results in the lowest overall formal charge.
The existence of this unpaired electron significantly impacts the molecule’s chemical behavior, making free radicals important in biological and atmospheric chemistry. Due to the instability of the lone electron, these species are extremely reactive and short-lived, constantly seeking another electron to achieve a stable, paired configuration. This high reactivity allows them to act as potent intermediates in chemical reactions, readily bonding with other molecules to neutralize the unpaired electron.