Molecular behavior begins with its electrons, specifically those found in the outermost shell, known as valence electrons. These are the electrons an atom uses to form chemical bonds with other atoms, determining the resulting molecular structure and reactivity. Sulfur dioxide (\(\text{SO}_2\)) is a colorless gas, a common atmospheric pollutant that forms during the burning of fossil fuels containing sulfur. Understanding the molecular properties of this compound requires first establishing the total number of its valence electrons. This fundamental count dictates how the atoms will connect and arrange themselves in three-dimensional space.
Calculating the Total Valence Electrons in Sulfur Dioxide
To determine the total number of valence electrons in the sulfur dioxide molecule, we must first look at the positions of the constituent atoms on the periodic table. Both Sulfur (S) and Oxygen (O) belong to Group 16, meaning that each atom possesses six valence electrons. The \(\text{SO}_2\) molecule contains one sulfur atom and two oxygen atoms. The single sulfur atom contributes six valence electrons. Since there are two oxygen atoms, they collectively contribute twelve electrons (\(2 \times 6\)). Sulfur dioxide is found to have a total of eighteen valence electrons available for bonding and lone pairs. This total count of eighteen electrons must be precisely distributed to create a stable compound that satisfies the octet rule for each atom where possible.
Drawing the Lewis Structure and Formal Charges
The next step is to arrange the eighteen valence electrons around the atoms to form a stable Lewis structure, beginning by placing the least electronegative atom, Sulfur, at the center. Initial single bonds are drawn connecting the central sulfur atom to the two terminal oxygen atoms, utilizing four of the eighteen total electrons. The remaining fourteen electrons are then distributed, first as lone pairs to satisfy the octets of the terminal oxygen atoms.
At this point, the two oxygen atoms each have a complete octet, but the central sulfur atom only possesses four electrons from the single bonds and has two remaining electrons. These final two electrons must be placed on the sulfur atom as a single lone pair. However, in this arrangement, the sulfur atom still only has six electrons, which is an incomplete octet and results in unfavorable formal charges.
To rectify the incomplete octet and minimize the formal charges, a lone pair from one of the oxygen atoms must be converted into a double bond with the sulfur atom. This action creates one sulfur-oxygen double bond and one sulfur-oxygen single bond, satisfying the octet for all three atoms. Because the double bond can form with either of the oxygen atoms, \(\text{SO}_2\) exhibits two equivalent resonance structures. The most accurate single Lewis representation minimizes all formal charges by having sulfur expand its octet, forming a double bond with both oxygen atoms.
Understanding the Molecular Shape and Polarity
Once the Lewis structure is established, the three-dimensional shape of the molecule can be determined using the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory states that electron domains around the central atom will orient themselves as far apart as possible to minimize repulsion. In sulfur dioxide, the central sulfur atom has three electron domains: two domains from the bonding pairs and one domain from the lone pair of electrons.
These three electron domains attempt to occupy the corners of a flat triangle, giving the molecule an electron-domain geometry described as trigonal planar. However, the molecular geometry, which describes only the position of the atoms, is different because the lone pair is not an atom. The lone pair of electrons exerts a greater repulsive force than the bonding pairs, pushing the two oxygen atoms closer together. This repulsion results in a bent or V-shaped molecular geometry with a bond angle of approximately \(119^\circ\).
The molecule’s polarity is a direct consequence of this asymmetrical bent shape and the difference in electronegativity between sulfur and oxygen. Oxygen is more electronegative than sulfur, meaning it pulls the shared electrons in the bonds closer to itself, creating a bond dipole moment for each sulfur-oxygen bond. Because the molecule is bent, the two individual bond dipoles do not cancel each other out, as they would in a linear molecule like \(\text{CO}_2\). Instead, they combine to produce a net dipole moment, making the entire sulfur dioxide molecule polar.