Chemistry
Hund's Rule
Quick fact
Hund's rule is named after German physicist Friedrich Hund, who formulated it in 1925 based on spectroscopic observations of atoms.
Why this is interesting
Have you ever wondered why some atoms are naturally magnetic? It all comes down to a simple rule about how electrons prefer to fill their orbitals: they'd rather be alone than share a room.
Read the full explanation
Understanding Hund's Rule
Imagine a multilevel parking garage where each level has several parking spots (orbitals). Each spot can hold at most two cars (electrons), but cars don't like to park next to another car unless all other spots on that level are already taken. Similarly, electrons in a subshell (like the 2p orbitals) will first occupy each orbital singly, with their spins aligned (parallel), before any orbital gets a second electron with opposite spin. This behavior is called 'maximum multiplicity' because it maximizes the number of unpaired electrons. For example, carbon has two electrons in the 2p subshell: they go into two different p orbitals with parallel spins, not into the same orbital paired up.
A deeper explanation
Hund's rule arises from two quantum mechanical factors: electron-electron repulsion and the exchange interaction. When electrons occupy separate orbitals, they are farther apart on average, reducing Coulomb repulsion energy. Additionally, parallel spins allow a quantum mechanical 'exchange' effect that lowers the total energy of the atom. This exchange energy is negative and only occurs when unpaired electrons have parallel spins. The combination of lower repulsion and exchange stabilization makes the high-spin configuration (more unpaired electrons) more stable. Hund's rule is essential for understanding why oxygen is paramagnetic (attracted to magnets) due to its two unpaired electrons, and why transition metals exhibit various magnetic and chemical properties. It also dictates the order of filling in electric configurations, guiding the prediction of atomic spectra and bonding behavior.