Hund’s Rule is one of the foundational principles used to determine electron configurations in IB Chemistry. While it may seem like a small detail, it underpins your understanding of atomic structure, periodic trends, and even magnetic properties. Examiners frequently test Hund’s Rule in Paper 1 multiple-choice questions and Paper 2 explanations involving electron configurations, orbital diagrams, and periodicity.
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Quick Start Checklist
Before diving deeper, make sure you know:
- Hund’s Rule applies to orbitals of equal energy (degenerate orbitals).
- Electrons occupy orbitals singly before pairing up.
- Electrons in singly occupied orbitals have parallel spins.
- The rule minimizes electron–electron repulsion.
- It helps predict magnetic properties and electron configurations.
These principles often appear when filling p, d, and f orbitals.
What Is Hund’s Rule?
Hund’s Rule states that electrons will occupy degenerate orbitals singly and with parallel spins before any orbital receives a second electron. Degenerate orbitals are orbitals of the same energy level—for example, the three p orbitals (px, py, pz) or the five d orbitals.
In simpler terms:
Electrons prefer to spread out evenly rather than pair up right away.
Why? Because electrons repel one another, and remaining unpaired reduces repulsion within an atom. This leads to a more stable configuration.
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Why Electrons Fill Orbitals Singly First
The main reason electrons avoid pairing is electron–electron repulsion. Two electrons in the same orbital must have opposite spins and remain close together, which increases repulsive forces. By spreading out among available orbitals first, electrons occupy positions that minimize energy.
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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