Hund's rule

Hund's Rule is a principle in quantum mechanics that guides the arrangement of electrons in the electron shells of atoms and molecules. According to this rule, electrons will fill degenerate orbitals (orbitals that have the same energy level) singly as much as possible, maintaining parallel spins, before pairing up in the same orbital. This behavior is attributed to the electrons' tendency to minimize electron-electron repulsions within the atom or molecule, which is a consequence of the Pauli exclusion principle and Coulomb's law. Hund's Rule is crucial for understanding the electronic configurations of atoms and the magnetic properties of atoms and molecules.

Overview

Hund's Rule can be summarized in three main points:

1. For a given electron shell, electrons occupy all degenerate orbitals singly before any orbital is doubly occupied.
2. Electrons in singly occupied orbitals have the same spin quantum number (parallel spins), which maximizes the total spin of the atom or molecule.
3. When electrons do pair in orbitals, they do so with opposite spins, to comply with the Pauli exclusion principle.

This rule is named after Friedrich Hund, a German physicist who first proposed it in 1925. Hund's Rule, together with the Aufbau principle and the Pauli exclusion principle, provides a foundational framework for determining the electron configurations of atoms in their ground states.

Application

Hund's Rule is applied in the building-up process of an atom's electron configuration. For example, in the case of the carbon atom, which has six electrons, the first two electrons fill the 1s orbital. The next two electrons occupy the 2s orbital. According to Hund's Rule, the remaining two electrons will occupy the 2p orbitals singly, with parallel spins, rather than pairing up in the same 2p orbital. This arrangement minimizes repulsion between the electrons and is energetically favorable.

Significance

Understanding Hund's Rule is essential for predicting the chemical and physical properties of elements. For instance, the rule helps explain why oxygen, with two unpaired electrons in its 2p orbitals, is paramagnetic (attracted to magnetic fields), while nitrogen, with three unpaired electrons in its 2p orbitals, has an even stronger paramagnetic effect. Hund's Rule also plays a critical role in the field of spectroscopy, as it affects the energy levels and thus the spectral lines of atoms and molecules.

Limitations

While Hund's Rule provides a general guideline for electron configurations, there are exceptions, especially for atoms in the d and f blocks of the periodic table. In these cases, electron-electron interactions and the energy differences between orbitals can lead to configurations that deviate from those predicted by Hund's Rule.

See Also

• Aufbau principle
• Pauli exclusion principle
• Electron configuration
• Quantum mechanics
• Spectroscopy