Electron Orbital Filling Principles

Explain Aufbau's principle, Hund's rule, and Pauli's exclusion principle and their application in filling electron orbitals.

In the realm of quantum mechanics, particularly when discussing atomic structure, the arrangement of electrons within an atom's orbitals is governed by three fundamental principles: Aufbau's principle, Hund's rule, and Pauli's exclusion principle. These principles serve as the foundation for understanding how electrons make their home in atoms, influencing chemical properties and reactivities.

Aufbau's principle, derived from the German word 'Aufbauen' meaning 'to build up', suggests that electrons occupy the lowest energy orbitals first before filling higher ones. This helps in predicting electron configurations in a systematic manner by following the order of increasing energy levels.

Hund’s rule provides insight into how electrons occupy degenerate orbitals, such as p, d, or f orbitals. It states that electrons will singly occupy all the degenerate orbitals before any orbital is doubly occupied. This minimizes repulsion among electrons and helps achieve a more stable configuration by maximizing the number of unpaired electrons with parallel spins.

The Pauli exclusion principle, one of the cornerstones of quantum mechanics proposed by Wolfgang Pauli, states that no two electrons can have the same set of quantum numbers within an atom. This means that each electron in an atom has a unique state, distinguished by its quantum numbers, effectively preventing any two electrons from being in the same place at the same time. Together, these principles shape our understanding of an atom’s electronic structure, guiding predictions in chemical behavior and interactions.