Hund’s Rule Explained for IB Chemistry

This is why, for example:

• Nitrogen’s configuration ends in 2p³ with one electron in each p orbital.
• Oxygen’s ends in 2p⁴ with one paired orbital and two singly occupied orbitals.

Understanding this pattern helps with questions on magnetic behavior, since unpaired electrons cause paramagnetism.

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Hund’s Rule and Magnetic Properties

Atoms with unpaired electrons exhibit paramagnetism (weak attraction to magnetic fields).
Atoms with all electrons paired exhibit diamagnetism (weak repulsion).

Examples:

• Oxygen (2p⁴) → 2 unpaired electrons → paramagnetic
• Neon (2p⁶) → all electrons paired → diamagnetic

Paper 2 often asks students to justify magnetic properties using orbital diagrams and Hund’s Rule, so practice drawing these diagrams with care.

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Hund’s Rule in Orbital Diagrams

Orbital diagrams visually represent electron placement. When applying Hund’s Rule:

• Fill each degenerate orbital with one electron before pairing.
• Use upward arrows (↑) for unpaired electrons.
• Use downward arrows (↓) to indicate paired electrons with opposite spin.

This helps you identify mistakes in electron configurations—a common source of lost marks on exams.

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Frequently Asked Questions

Why must electrons in singly occupied orbitals have parallel spins?

Parallel spins minimize electron repulsion and stabilize the atom. When electrons share the same spin, exchange energy increases, making the configuration lower in energy.

Does Hund’s Rule apply to all orbitals?

Hund’s Rule applies only to degenerate orbitals (orbitals of equal energy), such as p, d, and f subshells. It does not affect the order in which subshells fill—that’s governed by the Aufbau Principle.

How is Hund’s Rule tested in IB exams?

Expect questions involving orbital diagrams, electron configurations, periodic trends, and magnetic properties. Multiple-choice items often test recognition of correct or incorrect configurations.

Conclusion

Hund’s Rule ensures that electrons occupy orbitals in the most stable arrangement by minimizing repulsion and maximizing spin parallelism. Understanding this rule strengthens your grasp of atomic structure, periodic trends, and magnetic behavior—core ideas that support success throughout IB Chemistry.

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