Sulfur dioxide (\(\text{SO}_2\)) is a common gaseous compound frequently released into the atmosphere through volcanic activity and the combustion of fossil fuels. The bond in sulfur dioxide is best classified as polar covalent. This means the atoms share electrons, characteristic of covalent bonding, but they do so unequally, resulting in distinct electrical poles within the molecule. This unequal sharing dictates the physical properties and reactivity of the \(\text{SO}_2\) molecule.
The Chemical Foundation of Bonding
The type of bond formed between two atoms is determined by electronegativity, which is an atom’s ability to attract shared electrons toward itself in a chemical bond. Chemical bonds exist on a spectrum, with ionic bonds at one extreme and nonpolar covalent bonds at the other. Most compounds, including sulfur dioxide, exhibit covalent bonds where the electron sharing is unequal.
To determine the bond type in \(\text{SO}_2\), we examine the electronegativity values for sulfur (\(\text{S}\)) and oxygen (\(\text{O}\)). Oxygen is highly electronegative (3.5), while sulfur is less electronegative (2.5). The difference between these two values is 1.0, which confirms a polar covalent bond.
This difference in electron-pulling power means that when sulfur and oxygen atoms form a bond, the shared electrons spend more time closer to the more electronegative oxygen atom. The disproportionate distribution of the electron cloud creates a partial negative charge (\(\delta^-\)) on the oxygen atom and a partial positive charge (\(\delta^+\)) on the sulfur atom. This electrical separation within each individual sulfur-oxygen bond is the defining characteristic of its polarity.
Visualizing the Electron Arrangement and Resonance
The precise arrangement of electrons in sulfur dioxide is visualized using a Lewis structure, which shows how the atoms share their valence electrons. Sulfur and oxygen are both in Group 16, meaning each atom contributes six valence electrons, totaling eighteen for the \(\text{SO}_2\) molecule.
The central atom in \(\text{SO}_2\) is sulfur because it is the less electronegative element, and it is bonded to the two oxygen atoms. After forming a single bond between the sulfur and each oxygen atom, the remaining electrons are distributed to complete the octets. However, a single Lewis structure cannot fully represent the actual bonding in sulfur dioxide.
The complexity of the \(\text{SO}_2\) molecule requires the concept of resonance, which describes molecules where the true electron arrangement is an average of two or more valid Lewis structures. In one structure, the sulfur atom is double-bonded to one oxygen atom and single-bonded to the other, with a lone pair on the central sulfur atom. The second structure flips the position of the single and double bonds.
The actual bonding in sulfur dioxide is a hybrid of these two contributing structures, meaning the electrons are constantly delocalized across the entire structure. This results in the two sulfur-oxygen bonds being identical, neither a pure single bond nor a pure double bond. Instead, each bond is intermediate, possessing a bond order of approximately 1.5.
Molecular Geometry and Overall Polarity
The three-dimensional shape of the molecule is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory. This model posits that electron groups—whether they are bonding pairs or lone pairs—arrange themselves around the central atom to minimize repulsion. For sulfur dioxide, the central sulfur atom has three electron groups: two regions of shared electrons and one lone pair of electrons.
These three electron groups attempt to occupy the corners of a triangle, which would normally result in a trigonal planar electron geometry. However, the molecular geometry is different because the lone pair occupies more space than the bonding pairs, exerting a greater repulsive force on the two sulfur-oxygen bonds.
This repulsion pushes the two oxygen atoms closer together, resulting in a “bent” or “V-shaped” molecular geometry for \(\text{SO}_2\). The bond angle is slightly less than the ideal \(120^\circ\), measured to be approximately \(119^\circ\). This asymmetrical bent shape is the final factor in determining the molecule’s overall polarity.
Although the individual sulfur-oxygen bonds are polar, an overall molecular dipole moment only exists if the molecule is asymmetrical. Because the \(\text{SO}_2\) molecule is bent, the pull of the electrons towards the two oxygen atoms does not cancel out, resulting in a net dipole moment. This makes the sulfur dioxide molecule a polar molecule, influencing how it interacts with other substances.