Sulfur dioxide (\(\text{SO}_2\)) is a common gaseous molecule, often released from industrial processes, and is a significant air pollutant. When considering its fundamental molecular properties, sulfur dioxide is a polar molecule, meaning it possesses a separation of electrical charge across its structure. The polarity of \(\text{SO}_2\) results from the type of atoms involved and their specific arrangement in three-dimensional space.
The Polarity of the S-O Bond
Molecular polarity begins with bond polarity, which is determined by the electron-attracting ability of each atom, known as electronegativity. Electronegativity measures how strongly an atom pulls shared electrons toward itself within a chemical bond. In the \(\text{SO}_2\) molecule, the central sulfur atom is bonded to two oxygen atoms.
Oxygen is significantly more electronegative (3.44) than sulfur (2.58). This substantial difference causes the shared electrons in the sulfur-oxygen bonds to be pulled closer to the oxygen atoms. A partial negative charge (\(\delta-\)) develops on each oxygen atom, and a corresponding partial positive charge (\(\delta+\)) is left on the central sulfur atom.
This unequal sharing establishes a bond dipole moment within each individual S-O bond. The bond dipole moment is a vector pointing from the less electronegative sulfur toward the more electronegative oxygen. These directional forces confirm that the individual S-O bonds are polar.
How Molecular Geometry Shapes Polarity
While the bonds themselves are polar, the overall polarity of a molecule depends entirely on its three-dimensional shape, or molecular geometry. If the polar bonds are arranged symmetrically, their individual pulls can cancel each other out, resulting in a nonpolar molecule. The geometry of \(\text{SO}_2\) is dictated by the Valence Shell Electron Pair Repulsion (VSEPR) theory.
The central sulfur atom in \(\text{SO}_2\) has two bonding regions connected to the oxygen atoms and one non-bonding lone pair of electrons. The presence of this single lone pair prevents a symmetric structure from forming. The lone pair exerts a stronger repulsive force than the bonding pairs, pushing the two oxygen atoms closer together.
This repulsion results in a “bent” or “V-shaped” molecular geometry for \(\text{SO}_2\), with a bond angle of approximately \(119^\circ\). If sulfur dioxide were linear, like carbon dioxide (\(\text{CO}_2\)), the two bond dipoles would pull in opposite directions and perfectly cancel. Because the lone pair forces a bent shape, the molecule is asymmetrical. This asymmetrical arrangement means that the individual bond dipoles cannot be perfectly counteracted.
Why Sulfur Dioxide Is a Polar Molecule
The combination of polar bonds and an asymmetrical bent shape leads directly to the conclusion that sulfur dioxide is a polar molecule. Since the individual bond dipole moments are not pulling in opposite and equal directions, they do not cancel out. Instead, they add together vectorially to create a single, permanent net dipole moment for the entire \(\text{SO}_2\) molecule.
This net dipole moment points away from the central sulfur atom toward the region between the two more electronegative oxygen atoms. The existence of this net separation of charge across the molecule is the defining characteristic of a polar substance. Sulfur dioxide has a measurable dipole moment.
The polarity of sulfur dioxide has significant real-world implications. Polar molecules tend to dissolve readily in other polar substances, following the principle of “like dissolves like.” Because water is a highly polar solvent, the polar \(\text{SO}_2\) molecule is highly soluble in water. This solubility allows sulfur dioxide to dissolve in atmospheric moisture, where it chemically reacts to form sulfurous acid, a primary component of acid rain.