Hydrogen sulfide (\(\text{H}_2\text{S}\)) is a colorless gas recognizable by its distinct rotten-egg odor. Its three-dimensional arrangement of atoms, known as molecular geometry, is fundamental for understanding its chemical behavior. This geometry dictates many of the compound’s physical and chemical properties, including reactivity and how it interacts with other molecules.
Building the Foundation: The Lewis Structure of \(\text{H}_2\text{S}\)
The Lewis structure visualizes the distribution of valence electrons, which is the foundation for determining molecular shape. Sulfur, which belongs to Group 16 of the periodic table, contributes six valence electrons, while the two hydrogen atoms contribute one each, totaling eight valence electrons for the molecule. Sulfur is the central atom because hydrogen can only form a single bond. The central sulfur atom forms two single covalent bonds, using four of the total valence electrons.
The remaining four valence electrons are placed on the central sulfur atom as two non-bonding, or lone, pairs. This arrangement satisfies the octet rule for sulfur and the duet rule for each hydrogen atom. This configuration establishes a total of four electron domains around the central sulfur atom: two domains are bonding pairs, and two domains are lone pairs. These four electron domains are the structural elements that determine the molecule’s final three-dimensional shape.
Predicting the Three-Dimensional Shape
The final spatial arrangement of the atoms in \(\text{H}_2\text{S}\) is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory posits that all electron domains, whether bonding or non-bonding, repel each other and will arrange themselves in three-dimensional space to maximize the distance between them. Since the sulfur atom has four total electron domains, the electron domain geometry is tetrahedral, resulting in an idealized bond angle of \(109.5^{\circ}\).
The molecular geometry, however, only describes the positions of the atoms, not the lone pairs. Because the \(\text{H}_2\text{S}\) molecule has two bonding pairs and two lone pairs, its molecular geometry is described as bent or V-shaped. Lone pairs of electrons occupy more space and exert a greater repulsive force than bonding pairs, which are held between two atomic nuclei. This increased lone pair-lone pair and lone pair-bond pair repulsion pushes the two hydrogen atoms closer together than the ideal tetrahedral angle would suggest.
The actual measured bond angle for the \(\text{H}-\text{S}-\text{H}\) bond in hydrogen sulfide is approximately \(92^{\circ}\) to \(92.1^{\circ}\). This significant reduction from the \(109.5^{\circ}\) tetrahedral angle is a direct consequence of the strong repulsive forces exerted by the two lone pairs on the central sulfur atom.
How Molecular Shape Determines Polarity
The bent molecular shape of \(\text{H}_2\text{S}\) determines its polarity, which is an unequal distribution of electrical charge across the molecule. The covalent bond between hydrogen and sulfur is polar because sulfur is more electronegative than hydrogen. Electrons in the \(\text{S}-\text{H}\) bonds are pulled closer to the sulfur atom, creating a partial negative charge near the sulfur and partial positive charges on the hydrogen atoms.
If \(\text{H}_2\text{S}\) were linear, the two bond polarity vectors (dipole moments) would cancel each other out, resulting in a nonpolar molecule. However, the bent geometry ensures the molecule is asymmetrical, meaning the individual bond dipoles do not cancel. The resulting vector sum of the two \(\text{S}-\text{H}\) bond dipoles points toward the central sulfur atom, creating a net non-zero dipole moment. This net dipole moment, measured at about \(0.95\) Debye, confirms that \(\text{H}_2\text{S}\) is a polar molecule.