Hund's rule
/ HOONT /
Picture passengers boarding a half-empty bus with rows of double seats. Most people would rather take an empty row to themselves than squeeze in next to a stranger. Electrons filling a set of equal-energy orbitals behave the same way: given several seats of the same price, they spread out one to a seat before any of them doubles up.
Hund's rule states that when electrons fill orbitals of the same energy (a subshell, such as the three p orbitals or five d orbitals), they occupy them singly and with parallel spins as far as possible before pairing up. Only once every orbital in that set holds one electron do further electrons start sharing. This arrangement gives the lowest-energy, most stable ground state.
The reason is twofold: electrons that occupy separate orbitals stay farther apart and so repel each other less, and electrons with parallel spins gain a quantum-mechanical stabilisation called exchange energy. The caveat is that Hund's rule decides the ground state's spin arrangement; it does not override the higher principle that you fill lower-energy subshells first.
Carbon's two outer 2p electrons do not crowd into a single p orbital. Instead they take two different p orbitals, one each, with their spins aligned. As a result carbon has two unpaired electrons, which is why it is slightly attracted by a magnet (it is paramagnetic).
Spread out singly with parallel spins before pairing — the lowest-energy seating.
Hund's rule explains why half-filled subshells (like nitrogen's 2p³ or chromium's 3d⁵) are unusually stable: every orbital holds one parallel-spin electron, maximising exchange stabilisation.