Hydrogen sulfide (\(\text{H}_2\text{S}\)) is a colorless gas known for its strong, unpleasant odor reminiscent of rotten eggs. This chemical compound is formed when two hydrogen atoms bond with a single sulfur atom. The bond holding these atoms together is definitively covalent in nature, meaning the atoms are linked through a process of electron sharing. This shared-electron bonding dictates many of the physical and chemical properties of the \(\text{H}_2\text{S}\) molecule.
How Covalent and Ionic Bonds Differ
Chemical bonds are broadly categorized based on how atoms interact with their valence electrons. The two primary bond types are covalent and ionic, representing two extremes of electron interaction. Covalent bonds form when two atoms share one or more pairs of electrons, typically occurring between two nonmetal atoms. This sharing allows each atom to achieve a stable electron configuration.
By contrast, an ionic bond involves the complete transfer of one or more electrons from one atom to another. This transfer usually happens between a metal atom and a nonmetal atom. The loss or gain of electrons creates charged particles (ions) that are held together by a strong electrostatic attraction. The nature of the atoms involved provides a good initial indicator of the bond type.
Using Electronegativity to Classify Bonds
To precisely determine the type of bond, chemists use the concept of electronegativity. This is a measure of an atom’s tendency to attract a shared pair of electrons toward itself. The difference (\(\Delta\)EN) between the electronegativity values of the two bonded atoms is the key metric. A very small or zero difference suggests that electrons are shared equally, resulting in a nonpolar covalent bond.
As the electronegativity difference increases, the sharing of electrons becomes unequal, leading to a polar covalent bond. General guidelines suggest that a \(\Delta\)EN between approximately 0.4 and 1.7 signifies a polar covalent bond. If the difference is significantly greater than 1.7, a complete electron transfer has occurred, classifying the bond as ionic.
Determining the Bond Type in Hydrogen Sulfide
The specific bond type in hydrogen sulfide is determined by comparing the electronegativity of sulfur (2.58) to that of hydrogen (2.2). Calculating the difference in electronegativity (\(\Delta\)EN) between the two atoms yields a value of \(2.58 – 2.2 = 0.38\). Since both hydrogen and sulfur are nonmetals, the bond must be covalent, ruling out an ionic structure.
This \(\Delta\)EN value of 0.38 is low, but because it is not zero, the electrons are shared unequally. The bond is best described as polar covalent, albeit one with low polarity. Sulfur, being the more electronegative atom, exerts a slightly stronger pull on the shared electrons, creating a partial negative charge near the sulfur and a partial positive charge near each hydrogen atom.
The Polarity and Shape of the \(\text{H}_2\text{S}\) Molecule
The polarity of the \(\text{H}-\text{S}\) bonds and the molecular geometry are major contributors to the overall behavior of the molecule. The central sulfur atom in \(\text{H}_2\text{S}\) has two single bonds to hydrogen atoms and two unshared pairs of electrons (lone pairs).
According to the Valence Shell Electron Pair Repulsion (VSEPR) theory, these four electron groups arrange themselves to minimize repulsion. The two lone pairs occupy more space than the bonding pairs, distorting the shape from a perfect tetrahedron into a bent or V-shaped molecular geometry. The actual bond angle is approximately \(92^\circ\), smaller than the ideal \(109.5^\circ\) angle.
Because the molecule is bent and not linear, the slight dipoles from the two polar \(\text{H}-\text{S}\) bonds do not cancel each other out. This asymmetry results in \(\text{H}_2\text{S}\) having a net overall dipole moment, confirming that it is an overall polar molecule.