1. Hund's Rule and Electron-Electron Repulsion:
* Hund's Rule states that electrons will individually occupy orbitals within a subshell before pairing up in the same orbital. This is because electrons in different orbitals experience less repulsion from each other.
* When an orbital is completely filled, all electrons are paired, minimizing electron-electron repulsion. This minimizes the energy of the system, making it more stable.
2. Exchange Energy:
* When electrons have the same spin in different orbitals within a subshell, they can exchange positions. This exchange contributes to a stabilizing effect called exchange energy.
* In a completely filled orbital, all electrons are paired, maximizing the exchange energy, further contributing to stability.
3. Symmetry and Degeneracy:
* Completely filled orbitals have a higher degree of symmetry. This symmetry leads to a higher degeneracy of the orbitals, meaning they have the same energy level.
* This degeneracy, coupled with the minimized electron-electron repulsion, contributes to the stability of the system.
4. Shielding:
* Electrons in filled orbitals effectively shield the nucleus from the outer electrons. This means that the outer electrons experience a weaker attraction from the nucleus, making them less likely to be removed.
5. Lower Energy:
* The overall effect of these factors is that completely filled orbitals have lower energy than partially filled orbitals. This lower energy state corresponds to a more stable configuration.
In summary:
The combination of minimized electron-electron repulsion, maximized exchange energy, increased symmetry and degeneracy, and effective shielding all contribute to the enhanced stability of completely filled orbitals.
