The octet rule is a fundamental concept in chemistry, stating that atoms tend to bond to achieve eight valence electrons, similar to noble gases. This rule helps predict how atoms will combine and form molecules. However, while broadly useful, it is not universally applicable. Many elements and molecules exhibit stable configurations that deviate from this rule, showing that chemical stability can arise from diverse electron arrangements.
Elements That Form Incomplete Octets
Some elements form stable compounds with fewer than eight valence electrons around their central atom, a phenomenon known as an incomplete octet. This often occurs with elements that have a small number of valence electrons, as their small atomic size and the energetic costs involved make forming enough bonds for a full octet unfavorable.
Beryllium (Be) is a notable example, often forming compounds with only four valence electrons. In molecules like beryllium chloride (BeCl₂), beryllium shares its two valence electrons, forming two single covalent bonds.
Boron (B) similarly forms compounds with an incomplete octet, typically having six valence electrons. In boron trifluoride (BF₃), boron shares its three valence electrons with three fluorine atoms, resulting in six electrons around the central boron atom. This ‘sextet’ configuration is often more energetically favorable for boron. Aluminum, in the same group, also forms compounds with six valence electrons around its central atom.
Elements That Form Expanded Octets
Elements in the third period and beyond can accommodate more than eight valence electrons in their outermost shell, a concept known as an expanded octet. This ability stems from the availability of empty d-orbitals in their valence shells, which can participate in bonding, allowing the central atom to form more than four bonds and thus exceed the typical octet.
Phosphorus (P), located in Period 3, frequently demonstrates expanded octets. In phosphorus pentachloride (PCl₅), the central phosphorus atom forms five single bonds with chlorine atoms, resulting in ten valence electrons. The presence of accessible 3d orbitals allows phosphorus to utilize these for additional bonding.
Sulfur (S), also from Period 3, is another common example. In sulfur hexafluoride (SF₆), the sulfur atom bonds to six fluorine atoms, accumulating twelve valence electrons. The excitation of electrons into empty 3d orbitals enables sulfur to form these six covalent bonds. Even noble gases like Xenon (Xe) from Period 5 can form compounds with expanded octets, such as xenon tetrafluoride (XeF₄), highlighting that even traditionally inert elements can exceed the octet.
Molecules With Odd Electron Counts
Some molecules possess an odd total number of valence electrons, making it impossible for every atom to achieve a complete octet. These “odd-electron molecules” often contain at least one unpaired electron, classifying them as free radicals. This unpaired electron typically results in one atom having fewer than eight electrons in its valence shell.
Nitric oxide (NO) serves as a prime example. Nitrogen contributes five valence electrons and oxygen six, summing to eleven valence electrons. It is impossible to distribute these electrons for both atoms to achieve a full octet.
Nitrogen dioxide (NO₂) is another molecule with an odd electron count. It has seventeen valence electrons, with five from nitrogen and six from each of the two oxygen atoms. In its Lewis structure, the nitrogen atom typically ends up with only seven electrons. These odd-electron species are often highly reactive due to the instability of the unpaired electron.