Question

Draw the Lewis structure of ozone. Explain why the \(\mathrm{O}-\mathrm{O}\) bond \((1.28 \AA)\) is longer in ozone than in \(\mathrm{O}_{2}(1.21 \AA)\).

Short answer

The Lewis structure of ozone (O3) consists of two resonance structures with a central oxygen atom single-bonded to an outer oxygen atom on one side and double-bonded to an oxygen atom on the other side. Due to resonance, the bond order in O3 is 1.5, which is weaker than the O2 double bond with a bond order of 2. Consequently, the O-O bond in ozone is longer (1.28 Å) than the O-O bond in oxygen (1.21 Å).

Step-by-step solution

Determine the total number of valence electrons

Since ozone is composed of three oxygen atoms and each oxygen atom contributes six valence electrons, the total number of valence electrons in the molecule will be 3 x 6 = 18 valence electrons.

Determine the central atom

Usually, the least electronegative atom becomes the central atom. In the case of ozone, all three atoms are oxygen, so any of them can be the central atom. We will choose the leftmost oxygen atom to be the central atom for the purpose of our example.

Attach the other atoms to the central atom

Connect the central oxygen atom with the other two oxygen atoms by single bonds. This will use 2 x 2 = 4 valence electrons, leaving us with 14 valence electrons left.

Complete the octet for the outer atoms

In order to complete the octet for the outer oxygen atoms, add the remaining valence electrons as lone pairs. Each oxygen atom needs six electrons to complete its octet, so after allocating 6 electrons each to the outer oxygen atoms, there'll be 2 valence electrons left.

Place any remaining electrons on the central atom

The central oxygen atom currently has an incomplete octet. Assign the remaining 2 valence electrons as a lone pair on the central oxygen atom. Now, the central oxygen atom has an octet, but there is a need for a double bond in order to satisfy the octet for all oxygen atoms in the molecule.

Identify the resonance structures

To satisfy the octet requirement for all the oxygen atoms, we need to move a lone pair of electrons from one of the outer oxygen atoms to create a double bond with the central oxygen atom. There are two possibilities for this arrangement which leads to resonance structures in ozone. So the Lewis structure of ozone is represented by two resonance structures with a double bond between the central oxygen atom and one of the outer oxygen atoms.

Explain the bond length difference in O2 and O3

In O2, a double bond exists between the two oxygen atoms. In ozone (O3), the double bond is distributed between the two O-O bonds due to resonance. This causes the bond order in O3 to be between a single and a double bond (the bond order is 1.5) which is a weaker bond compared to the double bond in O2 (bond order of 2). Consequently, the O-O bond in ozone is longer (1.28 Å) than the O-O bond in oxygen (1.21 Å). To sum up, the Lewis structure of ozone consists of two resonance structures with a central oxygen atom single-bonded to an outer oxygen atom on one side and double-bonded to an oxygen atom on the other side. The bond length difference in O2 and O3 is due to the bond order in O3 being 1.5, leading to a weaker and longer bond in comparison to the double bond in O2 with a bond order of 2.

Key concepts

Valence Electrons

When we talk about valence electrons, we refer to the number of electrons that an atom possesses in its outermost shell and that can participate in the formation of chemical bonds. Oxygen, the building block of ozone (O3), has six valence electrons.

In creating a Lewis structure, the first step is to calculate the total number of valence electrons in the molecule. For ozone, with three oxygen atoms, this equates to \( 3 \times 6 = 18 \) valence electrons. Determining the total valence electrons is crucial since it lays the foundation for all subsequent steps in formulating the correct structure of the molecule, ensuring that every atom follows the octet rule wherever possible. The octet rule is the concept that atoms prefer to have eight electrons in their valence shell to be stable, resembling the electron arrangement of noble gases.

A good understanding of valence electrons not only aids in drawing Lewis structures but is also essential in predicting molecule geometry, bonding properties, and reactivity.

Resonance Structures

Resonance structures are a set of two or more Lewis Structures that collectively describe the electronic structure of a molecule, indicating that the actual distribution of the electrons is an average of the distribution indicated by the various structures. The need for resonance arises when one Lewis structure is not sufficient to represent a molecule accurately. Ozone presents an exemplary case of resonance.

The two resonance structures of ozone, indicated by double arrows, show that the double bond can exist between the central oxygen and either of the two outer oxygens. This implies that no single Lewis structure can represent the molecule completely; instead, the actual structure of ozone is a hybrid of these two resonance structures. These structures demonstrate that the negative charge and double bond are delocalized over the three oxygen atoms. Importantly, resonance structures also impact the molecule's stability and physical properties, such as bond lengths. The concept of resonance is vital in understanding the nature of chemical bonds in more complex molecules.

Bond Length and Bond Order

Bond length is the distance between the nuclei of two bonded atoms. When comparing the O-O bond lengths in ozone (O3) and molecular oxygen (O2), the disparity arises from the bond order concept, which is the number of chemical bonds between a pair of atoms. In O2, a bond order of 2 corresponds to a double bond, leading to a shorter bond length of \(1.21 \mathring{A}\).

In ozone, the bond order is not a whole number; because of the electron delocalization indicated by resonance structures, the bond order in ozone is 1.5—somewhere between a single and a double bond. This partial double-bond character causes ozone's O-O bonds to be weaker and longer (\(1.28 \mathring{A}\)) compared to the double bond in O2. Understanding bond length and bond order is essential for deciphering not just the stability of molecules but also their reactivity patterns and the mechanisms by which they react with other substances.