Lewis structures provide a visual representation of the valence electrons in a molecule, showing their arrangement as shared bonds or unshared lone pairs. This simple diagram is a powerful tool used in chemistry to predict how atoms connect, offering insight into a molecule’s overall shape and its potential chemical reactivity. The process of drawing a Lewis structure follows a systematic approach, beginning with a count of all available electrons. Hydrogen Sulfide (\(\text{H}_2\text{S}\)) is composed of two hydrogen atoms and one sulfur atom, serving as an excellent example for understanding these fundamental bonding principles.
Calculating Valence Electrons for \(\text{H}_2\text{S}\)
The first step in determining the Lewis structure is calculating the total number of valence electrons available for bonding within the \(\text{H}_2\text{S}\) molecule. Valence electrons are the electrons in the outermost shell of an atom, and their number is determined by the atom’s group number on the periodic table.
Sulfur (S) belongs to Group 16, contributing six valence electrons. Hydrogen (H) belongs to Group 1, contributing one valence electron per atom. Since \(\text{H}_2\text{S}\) indicates there are two hydrogen atoms, they collectively contribute two electrons. Adding these contributions—six from sulfur and two from hydrogen—yields a total of eight valence electrons. This total of eight electrons must appear in the final Lewis structure.
Arranging Atoms and Initial Bonding
The next step involves arranging the atoms to form the basic skeletal structure and drawing the initial bonds. Hydrogen atoms can never be central atoms because they can only form a single bond. Therefore, the single sulfur atom must be placed in the center, with the two hydrogen atoms positioned on the outside surrounding it.
A single covalent bond is represented by a line or two dots, signifying a shared pair of electrons. Drawing a single bond from the central sulfur atom to each of the two hydrogen atoms establishes the initial framework. These two single bonds utilize four of the total eight available valence electrons. This leaves four electrons remaining to be distributed as lone pairs to satisfy the stability requirements of the atoms.
Completing the Octet Rule with Lone Pairs
After establishing the initial bonds, the remaining valence electrons must be placed to satisfy the stability requirements of every atom in the molecule. The Octet Rule states that most atoms, including sulfur, must be surrounded by eight valence electrons to achieve a stable configuration. Hydrogen is an exception, requiring only two electrons (the Duet Rule).
In the \(\text{H}_2\text{S}\) structure, each hydrogen atom is already satisfied by the two electrons shared in its single bond with sulfur. The focus shifts to the central sulfur atom, which currently only has four electrons from the two single bonds. Since four electrons were used in the bonds, the four remaining valence electrons must be placed on the sulfur atom.
These four remaining electrons are placed on the sulfur atom as two non-bonding pairs, commonly called lone pairs. Placing these two lone pairs on the sulfur atom brings its total electron count to eight (four from shared bonds plus four from lone pairs), thereby fulfilling the Octet Rule. This arrangement satisfies the valence electron count and the stability rules for all atoms in the \(\text{H}_2\text{S}\) molecule.
The Final Lewis Structure and Molecular Shape
The completed Lewis structure for \(\text{H}_2\text{S}\) shows the central sulfur atom bonded to two hydrogen atoms with single covalent bonds, and the sulfur atom also possesses two lone pairs of electrons. This final diagram accurately represents the sharing and placement of the total eight valence electrons.
The presence of the two lone pairs on the central sulfur atom dictates the molecule’s three-dimensional arrangement. These non-bonding electron pairs exert a greater repulsive force than the bonding pairs, pushing the two hydrogen atoms closer together. This electronic repulsion results in a non-linear, angular shape for the molecule. Consequently, the molecular shape of Hydrogen Sulfide is described as “bent” or “V-shaped.”