Question

Compare the Lewis and molecular orbital treatments of the oxygen molecule.

Short answer

While both theories are integral to our understanding of molecular structure and behavior, Lewis' theory is much simpler and considers that electrons are localized between atoms; it unfortunately falls short when describing paramagnetism and other phenomena. Molecular Orbital Theory is more complex but also more comprehensive, explaining phenomena such as paramagnetism through the concept of electron sharing in molecular orbitals.

Step-by-step solution

Lewis Structure of Oxygen Molecule

First, let's talk about the Lewis structure of an oxygen molecule, \( O_2 \). Oxygen is in Group 6 on the periodic table, so an atom of oxygen has 6 electrons in its outer level. With two oxygen atoms, there are a total of 12 valence electrons. In \( O_2 \), you can see that each oxygen atom is connected to the other by a double bond which accounts for 4 of the 12 valence electrons. The remaining 8 electrons complete the octets of both the oxygen atoms. Hence, the Lewis structure shows us that oxygen molecule is a double bonded compound.

Molecular Orbital Theory of Oxygen Molecule

Let's now discuss the molecular orbital theory for \( O_2 \). According to this theory, when atomic orbitals combine, the number of molecular orbitals remains the same. For example, oxygen has two unpaired electrons in separate orbitals according to its molecular orbital diagram. Therefore, oxygen is paramagnetic. These properties cannot be revealed with Lewis structures which is considered as a limitation.

Comparing Both Treatments

Lewis structures provide a simple and easy-to-understand picture of bonding. However, some phenomena like the paramagnetism of \( O_2 \), cannot be explained with it. On the other hand, Molecular Orbital Theory, although mathematically intense, provides a more comprehensive description of molecular structure including the distribution of electrons in molecules.