Lewis structures illustrate the bonding and non-bonding valence electrons of atoms within a molecule. Sulfur, an element in Group 16, exhibits remarkable flexibility in its electron sharing, which directly influences the number of chemical bonds it can form. Its bonding capacity extends far beyond the simple configurations seen in many other elements. The number of bonds sulfur forms depends heavily on the molecule’s structure and the atoms it is bonded to.
Sulfur’s Standard Two-Bond Configuration
Sulfur’s most common bonding arrangement is determined by the octet rule. The sulfur atom has six valence electrons and typically needs two additional electrons to achieve stability. It accomplishes this by forming two covalent bonds, sharing one electron with each of two neighboring atoms.
This two-bond configuration is demonstrated in molecules such as hydrogen sulfide (\(\text{H}_2\text{S}\)). Sulfur forms a single bond with each of the two hydrogen atoms, using four valence electrons. The remaining four valence electrons exist as two non-bonding lone pairs, ensuring the sulfur atom is surrounded by eight electrons and satisfying the octet rule.
Sulfur primarily maintains this two-bond pattern in organic compounds like sulfides and thiols. The bonds may be two single bonds or occasionally one double bond, as seen in sulfur monoxide. This configuration represents the baseline bonding capacity for sulfur, mirroring the behavior of oxygen.
The Role of Expanded Valence Shells
The ability of sulfur to form more than two bonds stems from its position in the third period of the periodic table. Unlike second-period elements, sulfur has access to empty \(d\)-orbitals in its third principal quantum shell. These \(d\)-orbitals can be utilized to accommodate extra electrons when bonding demands greater sharing. This phenomenon, known as an expanded valence shell, allows the central sulfur atom to be surrounded by more than eight electrons.
The promotion of electrons into these \(d\)-orbitals unpairs some of sulfur’s lone pair electrons. Each newly unpaired electron can then form an additional covalent bond. This energetic promotion is favored when sulfur is bonded to highly electronegative atoms, such as oxygen or fluorine, which stabilize the expanded bonding state.
By utilizing its \(d\)-orbitals, sulfur can accommodate up to 12 valence electrons in its outer shell. This maximum corresponds to six shared pairs, meaning sulfur can form up to six covalent bonds. The existence of this expanded octet is why sulfur exhibits a wide range of bonding patterns.
Examples of Sulfur’s Variable Bonding
Sulfur’s capacity to access an expanded valence shell allows it to form molecules with four, five, or six bonds.
Four Bonds: Sulfur Tetrafluoride (\(\text{SF}_4\))
An example of sulfur forming four bonds is sulfur tetrafluoride (\(\text{SF}_4\)). In this molecule, sulfur bonds to four fluorine atoms, sharing electrons with each one. The sulfur atom in \(\text{SF}_4\) is surrounded by 10 valence electrons—four bonding pairs and one lone pair—which represents an expanded octet.
Six Bonds: Sulfur Hexafluoride (\(\text{SF}_6\))
A clear instance of sulfur forming its maximum of six bonds is sulfur hexafluoride (\(\text{SF}_6\)). Here, the sulfur atom is covalently bonded to six fluorine atoms, with each bond being a single shared pair of electrons. This results in the sulfur atom being surrounded by a total of 12 valence electrons, its theoretical maximum. The high electronegativity of the fluorine atoms facilitates this maximum expansion of the sulfur valence shell.
The Sulfate Ion (\(\text{SO}_4^{2-}\))
The sulfate ion (\(\text{SO}_4^{2-}\)) offers an example where the Lewis structure that minimizes formal charge shows six bonds. The central sulfur atom forms two single bonds and two double bonds with the four surrounding oxygen atoms. Although this requires 12 electrons around the sulfur atom, which violates the strict octet rule, it is the most chemically representative Lewis structure because it reduces the separation of charge across the molecule.