Hund's rule
Summary
Hund's rule says electrons fill degenerate orbitals (such as the three 2p orbitals) one at a time with parallel spins before any orbital gets a second electron. The most testable point is that it sets the number of unpaired electrons, which decides paramagnetism.
Detail
Meaning: Orbitals in the same subshell (the three 2p, five 3d, seven 4f) have equal energy, called degenerate. Hund's rule says the lowest-energy, ground-state arrangement puts one electron in each degenerate orbital, all with parallel spins, before any orbital receives a second electron. Pairing two electrons in one orbital forces them into the same region of space and raises repulsion, and parallel spins in separate orbitals lower the energy further.
On the MCAT: You are asked to draw an orbital diagram or count unpaired electrons. Nitrogen (2p³) has three unpaired electrons, oxygen (2p⁴) has two, and neon (2p⁶) has none. Unpaired electrons make an atom or ion paramagnetic (pulled into a magnetic field), while all paired electrons make it diamagnetic (weakly repelled). Transition metal ions work the same way: Fe³⁺ is [Ar]3d⁵ with five unpaired electrons.
Common confusion: Students mix this up with the Pauli exclusion principle. Pauli says an orbital holds at most two electrons, which must have opposite spins. Hund's rule decides how electrons spread across degenerate orbitals. It does not set the order in which subshells fill, which follows energy (the Aufbau principle).
Sources
- Chemistry: The Central Science
- Chemistry (Zumdahl)
- General Chemistry (Petrucci)
Written for the MCAT and checked before it was published. For study reference only. Spotted an error? Let us know.