Hund's Rule
Hund's rule of maximum multiplicity states that when electrons occupy degenerate (equal-energy) orbitals, one electron fills each orbital with parallel spins before any orbital receives a second electron. This arrangement minimises electron–electron repulsion because electrons in separate orbitals are farther apart, resulting in the lowest-energy (most stable) ground state. The rule also means that atoms with partially filled degenerate orbitals possess unpaired electrons, making them paramagnetic.
Worked Example
Problem
Draw the orbital diagram for nitrogen (N, Z = 7) and determine how many unpaired electrons it has.
Solution
Step 1 — Electron configuration of N: 1s² 2s² 2p³ Step 2 — Fill the three degenerate 2p orbitals using Hund's rule: 2p_x: ↑ (one electron, spin up) 2p_y: ↑ (one electron, spin up) 2p_z: ↑ (one electron, spin up) Step 3 — Each 2p orbital has exactly one electron; no pairing occurs. Step 4 — Count unpaired electrons: 3 (one in each 2p orbital).
Answer
Nitrogen has 3 unpaired electrons; it is paramagnetic.
Hund's Rule Applied to 2p Orbital Filling Across Period 2
| Element | Z | 2p Configuration | Unpaired Electrons | Magnetic Behaviour |
|---|---|---|---|---|
| Boron (B) | 5 | ↑ _ _ | 1 | Paramagnetic |
| Carbon (C) | 6 | ↑ ↑ _ | 2 | Paramagnetic |
| Nitrogen (N) | 7 | ↑ ↑ ↑ | 3 | Paramagnetic |
| Oxygen (O) | 8 | ↑↓ ↑ ↑ | 2 | Paramagnetic |
| Fluorine (F) | 9 | ↑↓ ↑↓ ↑ | 1 | Paramagnetic |
| Neon (Ne) | 10 | ↑↓ ↑↓ ↑↓ | 0 | Diamagnetic |
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Related Terms
Atomic Orbital
An atomic orbital is a mathematical function describing the wave-like behaviour and probable location of an electron in an atom, representing a region of space where there is a high probability (typically 90–95%) of finding the electron. Orbitals are characterised by three quantum numbers (n, l, mₗ) and have distinct shapes: s-orbitals are spherical, p-orbitals are dumbbell-shaped, and d- and f-orbitals have more complex geometries. Atomic orbitals form the basis for understanding electron configurations, chemical bonding, and molecular orbital theory.
Aufbau Principle
The Aufbau principle states that electrons fill atomic orbitals in order of increasing energy, beginning with the lowest-energy orbital available, before occupying higher-energy orbitals. The filling order follows the (n + l) rule: orbitals with lower (n + l) values are filled first; when two orbitals have the same (n + l) value, the one with the lower n is filled first. This principle, together with the Pauli exclusion principle and Hund's rule, allows chemists to predict the ground-state electron configuration of any element.
Electron Affinity
Electron affinity (EA) is the energy change that occurs when a neutral gaseous atom gains one electron to form a negative ion (anion). A negative EA value indicates an exothermic process — the atom releases energy and the anion is more stable than the separated atom and electron — which is the case for most halogens. Electron affinity generally increases (becomes more negative) across a period and decreases down a group, though there are notable exceptions such as the anomalously low EA of fluorine compared to chlorine due to electron–electron repulsion in fluorine's compact 2p orbitals.
Named after Friedrich Hund, a German physicist who formulated the rule in 1925 while studying atomic spectra. The full name is "Hund's rule of maximum multiplicity."