The octet rule explains how atoms bond by seeking eight electrons in their outermost, or valence, shell. Achieving this configuration provides stability, mimicking the electron arrangement of noble gases like Neon and Argon. While this concept accurately predicts bonding for many compounds, the complexity of chemical interactions means the rule is not universal and features several important exceptions.
Compounds with Fewer Than Eight Valence Electrons
Some elements form stable compounds even though the central atom is surrounded by fewer than eight valence electrons, a situation referred to as an incomplete octet. This violation occurs primarily with elements that possess very few valence electrons, making it difficult or impossible for them to achieve a full octet through typical covalent bonding. These electron-deficient molecules often seek out additional electrons to gain stability.
Boron is a prime example, having three valence electrons and typically forming three covalent bonds, as seen in Boron trifluoride (\(\text{BF}_3\)). In this molecule, the central Boron atom is associated with only six electrons. Similarly, Beryllium, with two valence electrons, forms compounds like Beryllium chloride (\(\text{BeCl}_2\)) where the central atom is surrounded by only four electrons.
For these elements, attempting to complete the octet by forming double bonds would result in unfavorable formal charges, especially on highly electronegative atoms like Fluorine. The incomplete octet is the more accurate depiction of the molecule’s reality because the resulting structure would otherwise be chemically unstable. Aluminum, located below Boron, also exhibits this behavior, often forming compounds like Aluminum chloride (\(\text{AlCl}_3\)) with only six valence electrons around the central atom.
Compounds with More Than Eight Valence Electrons
A class of exceptions involves atoms that can hold more than eight electrons in their valence shell, a condition known as an expanded octet. This phenomenon is strictly observed in elements found in the third period of the periodic table and beyond, such as Phosphorus, Sulfur, Chlorine, and Xenon. Elements in the second period, like Carbon, Nitrogen, and Oxygen, never exhibit expanded octets because they lack the necessary orbitals.
The ability to accommodate extra electrons is directly linked to the availability of low-lying, vacant \(d\)-orbitals in these heavier elements. While Period 2 elements only have \(2s\) and \(2p\) orbitals for bonding, elements in Period 3 and below have accessible \(3d\) orbitals that can participate in the bonding process. These empty \(d\)-orbitals allow the central atom to expand its valence shell to hold ten, twelve, or even more electrons.
Phosphorus pentachloride (\(\text{PCl}_5\)) illustrates this, where the central Phosphorus atom forms five covalent bonds, resulting in ten shared electrons. Sulfur hexafluoride (\(\text{SF}_6\)) provides an example with the Sulfur atom forming six bonds and having twelve valence electrons. Noble gases can also form stable compounds with expanded octets; for instance, the central Xenon atom in Xenon tetrafluoride (\(\text{XeF}_4\)) is surrounded by twelve electrons.
Molecules with an Odd Number of Electrons
A third exception involves molecules that possess an odd total number of valence electrons. Since electrons exist in pairs within stable molecules, an odd count means at least one electron must remain unpaired. These species are known as free radicals. Nitric Oxide (\(\text{NO}\)), with eleven valence electrons, and Nitrogen Dioxide (\(\text{NO}_2\)), with seventeen, are classic examples where it is impossible for every atom to be surrounded by exactly eight electrons.
The presence of this unpaired electron makes these molecules paramagnetic and extremely reactive. This single electron is readily available to form a new bond or react with other molecules, which is why free radicals are often short-lived. They seek to dimerize, or pair up, to achieve a more stable, even-electron configuration, such as when Nitrogen Dioxide forms Dinitrogen Tetroxide (\(\text{N}_2\text{O}_4\)), completing the octets for all atoms involved.