Sulfur’s capacity to form six bonds deviates significantly from the simple octet rule. This foundational principle suggests that atoms strive to achieve eight electrons in their outermost shell, which would limit sulfur to just two bonds. However, elements in the third period and beyond, including sulfur, can exceed this limit. Sulfur can accommodate up to 12 electrons in its valence shell, demonstrating a versatile bonding behavior that its lighter neighbors cannot match.
Sulfur’s Standard Electron Setup
Sulfur is positioned in Group 16 and Period 3, meaning a neutral atom has six electrons in its valence shell. These valence electrons occupy the $3s$ and $3p$ sublevels, resulting in the configuration $3s^23p^4$. This arrangement leaves two $3p$ orbitals with a single, unpaired electron.
Based on the simplest bonding model, sulfur would form two single covalent bonds to pair these electrons and achieve a stable octet. This pattern is seen in compounds like hydrogen sulfide ($\text{H}_2\text{S}$). Since sulfur can also form four or six bonds, it must employ a mechanism to make more valence electrons available for sharing.
Why Small Atoms Cannot Expand Their Bonding
The element directly above sulfur, oxygen, is strictly limited in its bonding capacity and cannot form six bonds. Oxygen atoms are located in Period 2, meaning their valence electrons are only found in the $2s$ and $2p$ orbitals. The second energy level contains only four orbitals—one $2s$ and three $2p$ orbitals—which hold a maximum of eight electrons total.
Period 2 atoms cannot accommodate more than eight electrons because there is no $2d$ sublevel. The next available empty orbitals are in the $3s$ sublevel, which is too high in energy to participate in bonding. This structural limitation forces elements like oxygen, nitrogen, and carbon to adhere strictly to the octet rule in most stable compounds.
The Mechanism of Expanded Bonding
The fundamental reason sulfur can form up to six bonds is its location in Period 3, which provides access to empty $3d$ orbitals. These $3d$ orbitals are energetically accessible, meaning they are close enough in energy to the filled $3s$ and $3p$ orbitals to participate in bonding. Although normally unoccupied in a neutral sulfur atom, they become involved when the atom forms more than two bonds.
To form six bonds, the sulfur atom uses promotion, absorbing energy to unpair electrons from the $3s$ and $3p$ orbitals and move them into the empty $3d$ orbitals. This excitation changes the arrangement from two unpaired electrons to six single, unpaired electrons. These six orbitals (one $3s$, three $3p$, and two $3d$) then hybridize to create six equivalent $sp^3d^2$ hybrid orbitals. This hybridization allows sulfur to accommodate 12 electrons in its valence shell, forming six single bonds with a stable, octahedral geometry.
Sulfur Compounds With Six Bonds
The theoretical capacity for six bonds is demonstrated in compounds like sulfur hexafluoride ($\text{SF}_6$), a highly stable and unreactive gas. In this molecule, the central sulfur atom forms six single covalent bonds with six fluorine atoms, placing 12 electrons around the sulfur center. The geometry of the $\text{SF}_6$ molecule is a perfect octahedron, providing maximum separation for the six bonding pairs.
The expanded bonding capacity is also demonstrated in the sulfate ion ($\text{SO}_4^{2-}$), a common component of substances like sulfuric acid. Although the actual bonding is complex, the most stable representation shows the sulfur atom forming six bonds to the surrounding oxygen atoms. This structure gives sulfur an expanded octet of 12 electrons, which helps minimize the formal charges on all atoms in the ion.