The number of lone pairs a sulfur atom possesses is not fixed; it depends entirely on the specific molecule or ion it is part of. A lone pair is a pair of valence electrons belonging to an atom that is not shared in a covalent bond. These non-bonding pairs substantially determine a molecule’s shape and chemical reactivity. To find the number of lone pairs, chemists must draw a Lewis structure. The total count is calculated based on sulfur’s initial valence electrons minus the electrons used for bonding.
The Foundation: Sulfur’s Valence Electrons and Ground State
To begin the calculation, establish the electron count for a neutral sulfur atom. Sulfur is element number 16 and belongs to Group 16 (chalcogens) on the periodic table. This placement dictates that a neutral sulfur atom possesses six valence electrons in its outermost shell.
The ground-state electron configuration of sulfur is \(1s^2 2s^2 2p^6 3s^2 3p^4\). These six valence electrons are distributed across the \(3s\) and \(3p\) orbitals. This pool of six electrons is available for forming bonds and existing as non-bonding pairs.
Chemical bonding is driven by the octet rule, where atoms seek a stable configuration of eight valence electrons. Sulfur typically achieves this by forming two covalent bonds to gain two additional electrons. The number of lone pairs is found by subtracting the bonding electrons from the total valence count and dividing the remainder by two.
The Standard Case: Sulfur Forming Two Bonds
The most common scenario where sulfur obeys the octet rule involves the formation of two single covalent bonds, a chemical situation analogous to that of oxygen in a water molecule. A prime example is hydrogen sulfide (\(\text{H}_2\text{S}\)), a compound where a central sulfur atom is bonded to two hydrogen atoms.
The calculation for sulfur in \(\text{H}_2\text{S}\) starts with six valence electrons. Forming a single bond with each of the two hydrogen atoms uses four electrons total. Subtracting these four bonding electrons from the initial six leaves two pairs of electrons not involved in bonding.
Therefore, the sulfur atom in a molecule like hydrogen sulfide has two lone pairs of electrons. These two lone pairs, along with the two bonding pairs, complete sulfur’s stable octet of eight electrons. The presence of these two lone pairs is crucial to the molecule’s three-dimensional arrangement.
Molecular Geometry of \(\text{H}_2\text{S}\)
According to Valence Shell Electron Pair Repulsion (VSEPR) theory, all electron pairs—both bonding and lone pairs—repel each other to maximize distance. Lone pairs take up more space than bonding pairs. The two lone pairs on the sulfur atom push the two hydrogen atoms closer together. This repulsion results in a bent or V-shaped molecular geometry for the \(\text{H}_2\text{S}\) molecule, with a bond angle of approximately 92.1 degrees.
When Sulfur Expands its Octet
Sulfur is in the third period, unlike second-period elements such as oxygen. This location gives it access to energetically available, unoccupied \(3d\) orbitals. This allows the sulfur atom to accommodate more than eight valence electrons, known as an expanded octet. When sulfur forms more than two bonds, it utilizes its valence electrons for bonding, reducing or eliminating its lone pairs.
Sulfur Dioxide (\(\text{SO}_2\))
The number of lone pairs decreases when electrons are used to form extra bonds. For instance, in sulfur dioxide (\(\text{SO}_2\)), the sulfur atom forms two double bonds with oxygen atoms. In its most common structure, it retains one lone pair of electrons. This single lone pair, combined with the bonding electrons, gives the sulfur a total of ten electrons, exceeding the standard octet.
Sulfur Hexafluoride (\(\text{SF}_6\))
In more extreme cases, the sulfur atom can use all six of its valence electrons to form six single bonds, resulting in a fully expanded octet of twelve electrons. A classic example of this is sulfur hexafluoride (\(\text{SF}_6\)), a compound where the sulfur atom is covalently bonded to six fluorine atoms. In this structure, all six of the sulfur’s valence electrons are used for bonding, leaving zero lone pairs on the central sulfur atom. The result is a highly symmetrical octahedral molecular geometry.