In traditional Lewis structures and molecular geometries, elements in the second period (period 2) of the periodic table, such as carbon, nitrogen, and oxygen, are generally depicted with an octet of electrons around them. This octet rule states that atoms tend to gain, lose, or share electrons in order to achieve a stable electron configuration with eight valence electrons, resembling the electron configuration of noble gases.

However, there are instances where elements can exceed the octet rule and accommodate more than eight electrons around them. This phenomenon is known as an expanded octet. Expanded octets are commonly observed in molecules or ions containing elements from the third period (period 3) of the periodic table and beyond, including elements such as sulfur, phosphorus, and chlorine.

The ability of these elements to form expanded octets arises from the availability of empty d orbitals in their valence shells. Elements in the third period and beyond have access to these d orbitals, which can participate in bonding and accommodate additional electrons beyond the octet.

Expanded octets are commonly observed in compounds involving these elements, particularly in molecules with central atoms surrounded by a larger number of bonded atoms or lone pairs. Examples of compounds with expanded octets include sulfur hexafluoride (SF6), phosphorus pentachloride (PCl5), and iodine heptafluoride (IF7).