Hund's Rule
Definition and meaning of Hund's Rule in chemistry.
Hund's rule states that electrons occupying degenerate orbitals (orbitals of equal energy within the same subshell) fill each orbital singly, with parallel spins, before any orbital receives a second electron.
In more detail
Placing electrons singly in separate orbitals with parallel spins minimizes electron-electron repulsion and maximizes the total spin multiplicity of the atom, which lowers its overall energy through a stabilization called exchange energy. Only after every orbital in the degenerate set holds one electron do additional electrons begin pairing up, with opposite spin, as required by the Pauli exclusion principle. This rule is essential for correctly predicting ground-state electron configurations, magnetic properties (paramagnetism versus diamagnetism), and atomic spectra.
Key facts
| Field | General Chemistry |
|---|---|
| Formal name | Hund's rule of maximum multiplicity |
| Proposed by | Friedrich Hund, 1925 |
| Applies to | Degenerate orbitals within the same subshell (e.g., the three p or five d orbitals) |
Nitrogen (Z = 7) has the configuration 1s2 2s2 2p3. Its three 2p electrons each occupy a separate 2p orbital (2px, 2py, 2pz) with parallel spins, rather than two electrons pairing in one 2p orbital while another remains empty.
Frequently asked questions
Why do electrons prefer to occupy separate orbitals with parallel spins?
Electrons with parallel spins in different orbitals experience less coulombic repulsion and gain extra stabilization (exchange energy) compared to two electrons paired in the same orbital, so the parallel-spin, singly-occupied arrangement has lower energy.
Does Hund's rule apply between orbitals of different subshells?
No. It applies only to a set of degenerate orbitals within the same subshell, such as the three 2p orbitals or the five 3d orbitals, not to orbitals of different energy like 2s and 2p.