paramagnetism of O2

Pour liquid oxygen between the poles of a strong magnet and it sticks, clinging to the field like a tiny blue puddle of iron filings. That is strange: oxygen is a simple gas of two identical atoms, and a magnet should not care about it. Yet O2 is genuinely magnetic, and explaining exactly why is the single most famous victory of molecular orbital theory.

Try drawing O2 the old way and you get a neat double bond, O=O, with every electron paired into bonds or lone pairs. Paired electrons cancel each other's tiny magnetism, so that picture predicts O2 should be non-magnetic. But it isn't. Molecular orbital theory tells a different story. After filling the lower orbitals, oxygen's last two valence electrons must go into two equal-energy pi-star antibonding orbitals. By Hund's rule they spread out, one in each, with the same spin — leaving two unpaired electrons. Unpaired electrons act like tiny magnets that do not cancel, so the molecule as a whole is paramagnetic, drawn to a magnetic field.

This matters far beyond a lecture demonstration. The two electrons sitting in antibonding orbitals also explain why O2's bond is a double bond (bond order two), weaker than nitrogen's triple bond. The unpaired-spin ground state, called triplet oxygen, is also why O2 is surprisingly sluggish to react despite being a strong oxidizer — many reactions with ordinary paired-electron molecules are spin-forbidden, which is part of why we and the wood around us do not simply burst into flame in air. One correctly read diagram explains the magnet trick, the bond strength, and the kinetics all at once.