Use the MO energy diagram in Figure 8.22b to describe the bonding in O2+, O2, and O2-. Which of the three is likely to be stable? What is the bond order of each? Which contain unpaired electrons?

Molecular Orbital (MO) Theory describes how atomic orbitals combine to form molecular orbitals, which can be occupied by electrons. In this theory, electrons are delocalized over the entire molecule, and the energy levels of these orbitals determine the stability and bonding characteristics of the molecule. Understanding the arrangement of electrons in these orbitals is crucial for predicting bond order and magnetic properties.

Bond order is a measure of the number of chemical bonds between a pair of atoms, calculated as the difference between the number of bonding and antibonding electrons divided by two. A higher bond order indicates a stronger bond and greater stability of the molecule. For example, in diatomic oxygen (O2), the bond order is 2, indicating a double bond, while in O2+ and O2-, the bond orders will differ due to the addition or removal of electrons.

Unpaired electrons are electrons that occupy an orbital alone rather than in pairs, contributing to the magnetic properties of a molecule. Molecules with unpaired electrons are often paramagnetic and can be less stable than those with all electrons paired. In the context of O2, O2+, and O2-, identifying the presence of unpaired electrons helps determine their stability and reactivity, with O2 being the most stable due to its paired electron configuration.