Final answer:
Oxygen’s unique electron configuration in diatomic molecules leads to a flip in the energy levels of pi and sigma antibonding orbitals due to electron-electron repulsion in its half-filled 2p subshell, explained by Hund's rule and molecular orbital theory.
Step-by-step explanation:
The energy of the pi orbitals in oxygen and other diatomic molecules is influenced by the electron configuration and the associated repulsive forces between electrons. In the case of oxygen with its electron configuration of 1s²2s²2p4, there is a particular consideration when it comes to these energy levels. When the 2p orbitals are more than half-filled, as they are in oxygen with four electrons in the 2p orbitals, the energy of the orbitals can be different compared to when they are less than half-filled. Specifically, this is why we observe a so-called 'flip' in energy levels where the pi* (antibonding) orbitals are at a lower energy than the sigma* (antibonding) orbitals for oxygen, which contrasts with other elements like nitrogen where the pi and sigma antibonding orbitals have the typical order.
This behavior relates to Hund's rule, where electrons in the same subshell (like 2p) will fill each orbital singly before pairing up. Oxygen's fourth electron in the 2p subshell pairs up, creating electron-electron repulsion which raises energy levels. This concept is extended in molecular orbital theory, explaining oxygen's unique properties such as its magnetic behavior and the structure of compounds like water, where hybridization of orbitals is required to explain the experimental bond angles.