Chemistry
Hund's Rule
Quick fact
Hund's rule was discovered by Friedrich Hund in 1927 and explains why atoms with unpaired electrons are attracted to a magnetic field (paramagnetic).
Why this is interesting
Imagine electrons as tiny, shy magnets. Why do they prefer to sit alone in separate rooms before sharing, even when there's plenty of space?
Read the full explanation
Understanding Hund's Rule
Electrons in an atom occupy orbitals—regions of space where they likely exist. Orbitals come in sets of equal energy, like three rooms (p orbitals) or five rooms (d orbitals). The Pauli exclusion principle says each room can hold at most two electrons, but they must have opposite spins. Hund's rule adds: when filling a set of equal-energy orbitals, electrons first occupy each orbital singly with the same spin direction (all 'up' or all 'down'), before any orbital gets a second electron. This minimizes repulsion because electrons spread out, and it maximizes the total spin, making the atom more stable.
A deeper explanation
The underlying reason is quantum mechanical: electrons repel each other because they have the same negative charge. Placing electrons in separate orbitals keeps them farther apart than pairing them in the same orbital. Additionally, parallel spins (same direction) allow the electrons to exchange positions without violating the Pauli principle, lowering energy through what physicists call 'exchange energy.' This exchange stabilization is strongest when spins are parallel and electrons are in different orbitals. The rule also predicts that an atom with the maximum number of unpaired electrons will have the lowest energy and highest spin multiplicity, which influences magnetic properties (paramagnetism). Hund's rule is essential for correctly writing electron configurations, especially for transition metals and rare earths, and for understanding why oxygen is paramagnetic and why iron is ferromagnetic.