In classical Lewis structures, molecular oxygen, O, is depicted as having a double bond between the two oxygen atoms. However, the molecular orbitals of O are actually more complex than Lewis structures seem to suggest. The highest occupied molecular orbital (HOMO) of O  is a pair of degenerate antibonding π orbitals, π* and π*, which are both singly occupied by spin unpaired electrons. These electrons are the cause of O being a triplet diradical in the ground state (indicated as O).
While many stable molecules’ HOMOs consist of bonding molecular orbitals and therefore require a moderate energy jump from bonding to antibonding to reach their first excited state, the antibonding nature of molecular oxygen’s HOMO allows for a lower energy gap between its ground state and first excited state. This makes excitation of O a less energetically restrictive process. In the first excited state of O, a 22 kcal/mol energy increase from the ground state, both electrons in the antibonding orbitals occupy a degenerate π* orbital, and oxygen is now in a singlet state (indicated as O). O is very reactive with a lifetime between 10-100µs.