Sulfur, a nonmetal element, holds a unique and flexible position in chemical bonding. Located in Group 16 of the periodic table, sulfur possesses six valence electrons in its outermost shell. Sulfur interacts with other atoms primarily through covalent bonds, where electron pairs are shared. The number of bonds sulfur can form is not fixed, but rather a variable that determines the structure and properties of the resulting molecule.
Sulfur’s Standard Bonding Limit (The Octet Rule)
In many compounds, sulfur adheres to the octet rule. This principle suggests that atoms bond to achieve eight electrons in their valence shell, similar to noble gases. Since a neutral sulfur atom starts with six valence electrons, it seeks to gain or share two more electrons to reach this stable count of eight.
This desire for two additional electrons dictates sulfur’s most frequent bonding pattern. In its ground state, sulfur has two unpaired electrons available for sharing. By forming a single covalent bond with two separate atoms, sulfur brings its total valence electron count to the stable eight.
A perfect example of this standard behavior is hydrogen sulfide (\(\text{H}_2\text{S}\)), where the sulfur atom forms a single bond with each of the two hydrogen atoms. This configuration satisfies the octet rule for sulfur while simultaneously satisfying the hydrogen atoms. In elemental sulfur, which exists as an eight-atom ring (\(\text{S}_8\)), each sulfur atom connects to two neighbors, maintaining the two-bond standard.
This two-bond structure is considered the default for sulfur, reflecting its most stable electronic arrangement when forming simple covalent compounds. The remaining two pairs of valence electrons on the sulfur atom that are not involved in bonding are known as lone pairs. These lone pairs influence the overall three-dimensional shape of the molecule.
Expanding the Bond Count: The Role of D-Orbitals
While the two-bond structure is common, sulfur can form four or even six covalent bonds, exceeding the standard octet rule. Sulfur is positioned in the third row of the periodic table, which allows for this expanded bonding capacity. Elements in the third period and beyond possess empty 3d-orbitals in their valence shell alongside the 3s and 3p orbitals.
These empty 3d-orbitals are energetically close to the filled 3s and 3p orbitals. When sulfur reacts, the energy released during bond formation can be used to promote, or excite, electrons into these empty d-orbitals. This excitation unlocks sulfur’s higher bonding potential.
In its ground state, sulfur has two unpaired electrons available for bonding. The first promotion of an electron from a lone pair to an empty 3d-orbital creates four unpaired electrons, allowing sulfur to form four covalent bonds. A second promotion, taking an electron from the 3s lone pair and moving it to a 3d-orbital, results in six unpaired electrons.
With six unpaired electrons, the sulfur atom is capable of forming six single covalent bonds. This process is often referred to as an “expanded octet” or hypervalency because the central sulfur atom is surrounded by more than eight valence electrons (twelve electrons for six bonds). The utilization of the d-orbitals allows sulfur to achieve various bonding configurations, significantly increasing its chemical versatility.
Common Molecules Showing Different Bond Numbers
The variable bonding capacity of sulfur is evident across a wide array of chemical compounds. The simplest and most common example of sulfur forming two bonds is hydrogen sulfide (\(\text{H}_2\text{S}\)), a toxic gas with a characteristic rotten egg smell. In this molecule, sulfur uses its two unpaired electrons to form a single bond with each hydrogen atom, achieving its stable octet.
When sulfur forms four bonds, the resulting compounds often contain double bonds or involve highly electronegative atoms. Sulfur dioxide (\(\text{SO}_2\)), a major air pollutant, features a sulfur atom with two double bonds to two oxygen atoms, resulting in a total of four bonds and ten valence electrons. Alternatively, in sulfur tetrafluoride (\(\text{SF}_4\)), the sulfur atom forms four single bonds with fluorine atoms, corresponding to a four-bond configuration.
The maximum number of bonds sulfur typically forms is six, demonstrated by compounds with the highest oxidation states. Sulfur hexafluoride (\(\text{SF}_6\)) is an inert, non-toxic gas where the sulfur atom is bonded to six fluorine atoms. This configuration arises from the maximum electron promotion into the d-orbitals, giving the sulfur atom six separate single bonds. Another example is sulfur trioxide (\(\text{SO}_3\)), where the sulfur atom forms six bonds through three double bonds with oxygen atoms.