Hund's rule: definition in chemistry
Hund's rule: Electrons fill orbitals of equal energy one at a time, with parallel spins, before any of them pair up.
Electrons in separate orbitals are farther apart and repel each other less, so spreading out across a sublevel gives a lower-energy arrangement than pairing up early. Hund's rule is used together with the aufbau principle and the Pauli exclusion principle to draw orbital diagrams. It explains why atoms such as nitrogen and oxygen have unpaired electrons, which make substances attracted to magnetic fields.
Examples
- Carbon (2p²): two electrons in separate p orbitals, both unpaired
- Nitrogen (2p³): one electron in each p orbital, three unpaired
- Oxygen (2p⁴): one pair and two unpaired electrons
- Iron (3d⁶): one pair and four unpaired electrons
On the periodic table
Across the p-block, the number of unpaired electrons rises from one in group 13 to three in group 15, then falls back to zero at the noble gases, just as Hund's rule predicts.