Lewis structures offer a simplified, two-dimensional visual model to represent the bonding and non-bonding electrons within a molecule. These diagrams are fundamental tools in chemistry, providing insight into how atoms share or transfer electrons to achieve stability. The arrangement of these electrons dictates a molecule’s chemical reactivity and physical properties. For the gas hydrogen sulfide (H₂S), the Lewis structure is the starting point for predicting its behavior and illustrating the basic principles of molecular formation.
Calculating Total Valence Electrons
The first step in constructing a Lewis structure is determining the total number of valence electrons available. Valence electrons are located in the outermost shell of an atom and are the only ones involved in chemical bond formation. This total count must be precisely accounted for in the final molecular diagram.
Hydrogen (H) is in Group 1, contributing one valence electron per atom. Since H₂S contains two hydrogen atoms, they collectively provide two electrons. Sulfur (S) is in Group 16, meaning it contributes six valence electrons.
Combining these contributions yields the total electron count for H₂S. Adding the two electrons from the hydrogen atoms to the six electrons from the sulfur atom results in a total of eight valence electrons. These eight electrons, or four electron pairs, must be distributed in the final Lewis structure.
Identifying the Central Atom and Initial Bonds
Once the total electron count is established, the next step is identifying the central atom and forming the initial skeletal structure. The central atom is typically the least electronegative element, though hydrogen is always positioned on the periphery because it can only form one bond. In H₂S, sulfur is less electronegative than hydrogen and takes the central position.
The two hydrogen atoms are connected to the central sulfur atom using single covalent bonds. Each single bond represents a shared pair of two electrons. Drawing these two initial bonds uses four of the total eight valence electrons, establishing the foundational structure before non-bonding electrons are considered.
Placing Lone Pairs and Verifying the Octet Rule
After placing the initial bonding electrons, the remaining valence electrons must be distributed as lone pairs. Subtracting the four bonding electrons from the total of eight valence electrons leaves four electrons yet to be placed. Hydrogen atoms are satisfied with the two electrons they already have from the single bonds, so no additional electrons are placed on them.
These remaining four electrons must be placed exclusively on the central sulfur atom. They are added as non-bonding electron pairs (lone pairs) to fill the available orbitals around the sulfur atom. This results in the sulfur atom possessing two lone pairs in addition to its two single bonds.
The final check confirms that the sulfur atom has achieved a stable electron configuration, known as a complete octet. The sulfur atom is surrounded by four electrons from the two single bonds and four electrons from the two lone pairs, totaling eight valence electrons. This arrangement of two bonds and two lone pairs completes the Lewis structure for H₂S.
The Molecular Geometry and Polarity of H₂S
The completed Lewis structure provides the information necessary to predict the molecule’s three-dimensional shape, determined by the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR posits that electron domains (bonding pairs and lone pairs) repel each other and arrange themselves in space to maximize the distance between them. The central sulfur atom in H₂S has four electron domains: two single bonds and two lone pairs.
These four domains initially arrange themselves in a tetrahedral electron geometry around the sulfur atom, ideally resulting in a bond angle of 109.5 degrees. However, lone pairs occupy more space and exert a greater repulsive force than bonding pairs, pushing the two hydrogen atoms closer together. The resulting three-dimensional structure is known as a “bent” or V-shaped molecular geometry.
The observed bond angle is approximately 92 degrees, a significant deviation from the ideal tetrahedral angle caused by the repulsion of the two lone pairs. This asymmetry is responsible for the molecule’s polarity. Sulfur has a higher electronegativity than hydrogen, pulling the bonding electrons closer to the sulfur atom and creating partial negative and positive charges across the molecule.
Because the molecule is bent, the individual bond dipole moments do not cancel each other out. This uneven distribution of electrical charge makes hydrogen sulfide a polar molecule. The molecular geometry and polarity contribute to the physical and chemical properties of H₂S.