Ozone, a molecule composed of three oxygen atoms (\(\text{O}_3\)), plays a significant role in atmospheric chemistry. Although its simple formula suggests a straightforward structure, the bonding is complex. The conventional Lewis structure model, which uses lines to represent fixed electron pairs, fails to accurately depict the electron distribution. This limitation necessitates the use of resonance, which describes molecules whose true structure is an average of multiple hypothetical forms.
Counting Valence Electrons and Connectivity
The first step in determining the structure is counting the total valence electrons. Each oxygen atom has six valence electrons, resulting in a total of eighteen (\(3 \times 6 = 18\)) for ozone.
The three atoms must be connected in an \(\text{O-O-O}\) arrangement, with one oxygen serving as the central atom. The initial structure uses two pairs of electrons to form single bonds between the central oxygen and the two outer atoms, accounting for four of the eighteen total electrons.
The remaining fourteen valence electrons are distributed to satisfy the octet rule. Placing three lone pairs on each terminal oxygen and one lone pair on the central atom uses all fourteen electrons. However, the central oxygen atom only possesses six electrons, two short of a complete octet. To rectify this, a lone pair from a terminal oxygen atom must be converted into a shared pair, forming a double bond with the central atom.
Deriving the Two Resonance Structures of Ozone
The process of moving a lone pair to form a double bond leads to the molecule’s first complete Lewis structure, featuring one single bond and one double bond. Since the lone pair used to form the double bond could have come from either terminal oxygen atom, two distinct and equally valid Lewis structures can be drawn for ozone.
In the first structure, the double bond exists between the central oxygen and the left terminal oxygen, while a single bond connects the central atom to the right terminal oxygen. The second structure is the mirror image, with the double bond placed on the right side and the single bond on the left.
The two structures are conventionally linked by a double-headed arrow, the chemical notation used to denote resonance. The movement of electrons required to convert one structure into the other involves shifting the multiple bond. This electron delocalization defines the two main resonance contributors for the ozone molecule. Since both structures are chemically equivalent and obey the octet rule, they are considered the only two significant resonance structures for ozone.
Evaluating Structure Stability Using Formal Charge
To confirm that these two structures are chemically plausible, the concept of formal charge is used to evaluate the electron distribution within each representation. Formal charge is a theoretical charge assigned to an atom in a molecule, calculated by comparing the number of valence electrons an isolated atom has with the number of electrons assigned to it in the Lewis structure.
In both resonance structures of ozone, the central oxygen atom is bonded to three pairs of electrons and possesses one lone pair, resulting in a formal charge of \(+1\). The terminal oxygen atom involved in the double bond has two lone pairs, giving it a formal charge of \(0\). The terminal oxygen atom involved in the single bond has three lone pairs, yielding a formal charge of \(-1\).
The two resonance structures are therefore equivalent, each exhibiting the same pattern of formal charges: one atom is \(-1\), one is \(+1\), and one is \(0\). Although a structure with zero formal charges on all atoms is generally preferred, this is not possible for ozone while maintaining the octet rule. The overall charge of the molecule is neutral, as the sum of the formal charges (\(+1 + 0 + (-1)\)) equals zero. The equivalence of the two forms confirms that they are the two primary contributors to the molecule’s true nature.
Understanding the Resonance Hybrid
The two Lewis structures are theoretical models and do not represent the actual physical state of the ozone molecule. The true structure is a blend of these two resonance forms, known as the resonance hybrid. In the hybrid, the electrons forming the multiple bond are delocalized over all three oxygen atoms.
This delocalization results in both oxygen-oxygen bonds being identical, with a bond order of approximately \(1.5\). This means each bond is stronger than a single bond but weaker than a double bond. Experimental measurements confirm this, showing that the two bond lengths are equal, measuring approximately \(127.8\) picometers \((\text{pm})\).
This measured length is intermediate between the typical length of an oxygen-oxygen single bond (\(147\text{ pm}\)) and a double bond (\(121\text{ pm}\)). The central oxygen atom carries a partial positive charge, while the two terminal oxygen atoms equally share the negative charge, each possessing a partial charge of \(-0.5\).