The principle that an electron only occupies one orbital at a time is rooted in the Pauli exclusion principle, which is a fundamental principle of quantum mechanics. The Pauli exclusion principle states that no two identical fermions (particles with half-integer spin, including electrons) can occupy the same quantum state simultaneously.
In the context of atoms, the Pauli exclusion principle implies that within a given atom, no two electrons can have the same set of quantum numbers. The quantum numbers describe various properties of an electron, including its energy level, orbital shape, orientation in space, and spin.
Each electron within an atom is characterized by a unique combination of these quantum numbers, and as a result, no two electrons can occupy the same orbital. This is often referred to as the "one-electron-per-orbital" rule.
The principle is a consequence of the wave-like behavior of electrons and the nature of quantum mechanical wave functions. The wave function describes the probability distribution of finding an electron in a particular region of space. The Pauli exclusion principle arises from the requirement that the wave function must be antisymmetric for identical fermions.
In practical terms, this means that electrons within an atom distribute themselves among different orbitals according to specific rules. For example, in the filling of atomic orbitals, the electrons occupy the lowest energy orbitals first, following the Aufbau principle. They fill orbitals in a way that maximizes the total spin and minimizes electron-electron repulsion.
The one-electron-per-orbital rule and the Pauli exclusion principle play a crucial role in determining the electronic structure and chemical behavior of atoms. They contribute to the stability and organization of matter by preventing the collapse of electrons into the same quantum state, leading to the formation of distinct energy levels and electron configurations within atoms.