Sulfur trioxide (\(text{SO}_3\)), consisting of one sulfur atom and three oxygen atoms, exhibits resonance structures. This molecule cannot be accurately represented by a single, static structural drawing because the bonding electrons are not confined between just two atoms. Resonance is a concept used when a single Lewis structure is insufficient to show the true nature of the electron distribution. The actual structure is a blend of several possible arrangements, each contributing to the molecule’s overall characteristics.
Understanding Electron Delocalization
The concept of resonance illustrates electron delocalization, which describes how electrons are shared among multiple atoms rather than being fixed in place. Delocalization means the electrons’ probability distribution is spread out over the entire molecule simultaneously, not oscillating between defined structures. This is similar to thinking of a blended smoothie, which is a single, uniform mixture, rather than a rapid alternation between the discrete fruits used to make it.
The electrons involved in the double bond are mobile and spread across the sulfur atom and all three oxygen atoms. This continuous sharing of electron density is a stabilizing influence, effectively lowering the molecule’s potential energy. Molecules that distribute their electrons over a larger area through resonance are inherently more stable. The theoretical Lewis structures we draw are contributors to the molecule’s actual state, which is called the resonance hybrid.
Deriving the Base Lewis Structure of Sulfur Trioxide
Constructing the Lewis structure for sulfur trioxide begins with calculating the total number of valence electrons available for bonding. Both sulfur and oxygen are in Group 16, meaning each atom contributes six valence electrons, for a total of 24 electrons (6 from sulfur + 3 x 6 from oxygen). The less electronegative sulfur atom is placed at the center, bonded to the three oxygen atoms.
The initial arrangement involves forming single bonds between the central sulfur and each of the three surrounding oxygen atoms, using six of the 24 available electrons. Placing the remaining 18 electrons as lone pairs around the oxygen atoms completes the octet for all three oxygen atoms. This preliminary structure is chemically unfavorable because it results in substantial formal charges on the atoms. The central sulfur atom would carry a +3 formal charge, and each oxygen atom would have a -1 formal charge, indicating a highly unstable arrangement.
To achieve a more stable configuration, the formal charges must be minimized, a process that often involves forming double bonds. Shifting a lone pair from one oxygen atom to form a double bond reduces the formal charge on that oxygen to zero and the sulfur charge from +3 to +2. A further reduction is achieved by converting another single bond to a double bond. The resulting configuration—one double bond and two single bonds—minimizes formal charges (+1 on sulfur, 0 on one oxygen, and -1 on the other two). This structure serves as the most representative Lewis structure for drawing the resonance forms.
The Three Equivalent Resonance Forms
The structure with one sulfur-oxygen double bond and two sulfur-oxygen single bonds is one of three possible equivalent structures. Since all three oxygen atoms in sulfur trioxide are chemically identical, the double bond has an equal probability of forming with any one of them. This results in three distinct, equally valid resonance forms, which are conventionally shown with a double-headed arrow connecting them.
The \(text{SO}_3\) molecule does not rapidly switch between these three structures. Instead, the true electron arrangement is a hybrid of all three contributing forms. This resonance hybrid suggests that the electron density associated with the double bond is delocalized and equally shared among all three sulfur-oxygen bonds.
Consequently, the actual \(text{SO}_3\) molecule has three bonds of equal length, each of which is an intermediate between a pure single bond and a pure double bond. The true bond order for each sulfur-oxygen linkage is calculated as the total number of bonding electron pairs (four) divided by the number of bonding regions (three), yielding a bond order of \(1 frac{1}{3}\). The measured bond length in \(text{SO}_3\) is approximately 1.42 Å, confirming the intermediate nature of the bonds.
Molecular Geometry and Polarity
The physical shape of the \(text{SO}_3\) molecule is a direct consequence of the bonding and electron delocalization. The central sulfur atom is surrounded by three bonding regions, one for each oxygen atom, and it does not possess any non-bonding lone pairs. This arrangement dictates a Trigonal Planar geometry, positioning the three oxygen atoms as far apart as possible in the same plane.
The bond angle between any two oxygen atoms is exactly 120°, creating a highly symmetrical structure. Although individual sulfur-oxygen bonds are polar due to the difference in electronegativity, the molecule itself is nonpolar. Because the three bond dipoles are of equal magnitude and arranged symmetrically at 120°, their effects cancel out completely, resulting in a net dipole moment of zero.