Hund's Rule Explained- Khan Academy Study Guide
What Is Hund's Rule?
Hund's Rule states that electrons fill orbitals of the same energy (degenerate orbitals) one at a time before pairing up. Each electron enters its own orbital with parallel spins.
That's the core definition. Now let's break down why this matters and how to actually use it.
Why Electrons Act This Way
Electrons carry negative charge. When two electrons occupy the same orbital, they repel each other. That's why electrons prefer to take empty orbitals first—they're avoiding that repulsion.
Parallel spins matter because of quantum mechanics. Electrons with parallel spins can exist in the same region of space more easily than electrons with opposite spins. It's called spin degeneracy, and it's the reason Hund's Rule exists.
In plain terms: electrons are lazy. They'll spread out before they double up.
Hund's Rule vs. The Other Electron Configuration Rules
You can't understand Hund's Rule in isolation. There are three main rules governing how electrons arrange themselves:
- Aufbau Principle: Fill lowest energy levels first (1s, then 2s, then 2p, etc.)
- Pauli Exclusion Principle: Each orbital holds maximum 2 electrons with opposite spins
- Hund's Rule: Fill degenerate orbitals singly before pairing
These rules work together. You use Aufbau to determine which orbitals to fill, then Hund's and Pauli to determine how to fill them.
The Practical Part: How to Apply Hund's Rule
Here's the step-by-step process:
Step 1: Identify Degenerate Orbitals
Degenerate orbitals are orbitals with the same energy level. The p subshell has 3 degenerate orbitals (px, py, pz). The d subshell has 5. The f subshell has 7.
Step 2: Fill Each Orbital Singly First
Put one electron in each orbital before adding a second to any of them. All first electrons get the same spin direction.
Step 3: Pair Up Last
Only after every orbital has one electron do you start adding the second electrons. These second electrons must have opposite spins.
Real Examples
Carbon (C) - Atomic Number 6
Carbon has 6 electrons. The configuration is 1s² 2s² 2p².
The 2p subshell holds the last 2 electrons. Following Hund's Rule:
Each 2p orbital gets one electron first. Then the second electrons go in—but there's only 2 electrons and 3 orbitals, so one orbital stays empty.
Visual representation:
- 2px: ↑
- 2py: ↑
- 2pz: (empty)
Not:
- 2px: ↑↓
- 2py: (empty)
- 2pz: (empty)
The second arrangement violates Hund's Rule and is wrong.
Nitrogen (N) - Atomic Number 7
Nitrogen has 7 electrons. Configuration: 1s² 2s² 2p³.
The 2p³ means 3 electrons in the p subshell. With Hund's Rule, each gets its own orbital:
- 2px: ↑
- 2py: ↑
- 2pz: ↑
All spins parallel. Maximum unpaired electrons. This is the most stable arrangement.
Oxygen (O) - Atomic Number 8
Oxygen has 8 electrons. Configuration: 1s² 2s² 2p⁴.
The 2p⁴ means 4 electrons in 3 orbitals. Fill singly first, then pair:
- 2px: ↑↓
- 2py: ↑
- 2pz: ↑
One orbital holds a paired set. The other two hold single electrons.
Orbital Diagrams: The Visual Method
Most chemistry courses require you to draw orbital diagrams. Here's the notation:
- Each box = one orbital
- Each arrow = one electron
- Up arrow = +½ spin
- Down arrow = -½ spin
For nitrogen's 2p subshell, the correct diagram looks like three boxes with up arrows in each. The wrong diagram shows two boxes with paired arrows and one empty box.
Quick Reference Table
| Element | 2p Electrons | Correct Arrangement | Unpaired Electrons |
|---|---|---|---|
| Carbon (6) | 2 | ↑ ↑ _ | 2 |
| Nitrogen (7) | 3 | ↑ ↑ ↑ | 3 |
| Oxygen (8) | 4 | ↑↓ ↑ ↑ | 2 |
| Fluorine (9) | 5 | ↑↓ ↑↓ ↑ | 1 |
| Neon (10) | 6 | ↑↓ ↑↓ ↑↓ | 0 |
Common Mistakes Students Make
Pairing too early: Putting two electrons in one orbital before filling all available orbitals with one electron each. This is the most frequent error.
Forgetting spin direction: All first electrons in degenerate orbitals must have parallel spins. You can't have one up and one down when filling singly.
Confusing Hund's Rule with Pauli: Pauli says two electrons in the same orbital must have opposite spins. Hund's says electrons in different orbitals of the same energy prefer parallel spins. These rules apply at different stages.
Why This Matters Beyond the Exam
Hund's Rule explains molecular bonding behavior. Unpaired electrons make atoms reactive. That's why nitrogen has 3 unpaired electrons and is so reactive—it wants to form bonds to pair those electrons.
Magnetic properties of elements depend on unpaired electrons too. Substances with unpaired electrons are paramagnetic. Those with all electrons paired are diamagnetic.
Oxygen is paramagnetic—it's attracted to magnetic fields. This is direct evidence that O₂ has unpaired electrons, which aligns with Hund's Rule predictions.
Getting Started: Your Study Checklist
- Memorize the definition: one electron per orbital first, parallel spins
- Practice drawing orbital diagrams for elements 1-20
- Identify degenerate orbital sets (p, d, f subshells)
- Check your diagrams: first pass fills each box with up arrows only
- Second pass adds down arrows starting from the left
- Count unpaired electrons and verify against the element's properties
Work through carbon through neon until the pattern clicks. It's a mechanical process—once you see the rhythm, you'll get every question right.