Why do electrons fill lower energy orbitals first?

Short Answer

 

Electrons fill lower energy orbitals first because atoms tend to achieve the lowest possible energy, which is the most stable arrangement. This principle is known as the Aufbau principle.

By occupying the lowest energy orbitals first, electrons minimize repulsion and maximize stability. This ordered filling determines the electron configuration of atoms, affects chemical reactivity, and explains the structure of the periodic table.

Detailed Explanation :

Reason for Filling Lower Energy Orbitals First

In atoms, electrons are negatively charged and are attracted to the positively charged nucleus. The energy of an electron depends on its distance from the nucleus and the type of orbital it occupies. Orbitals closer to the nucleus generally have lower energy, while those farther away have higher energy.

Principle of Minimum Energy

  • The Aufbau principle states that electrons occupy orbitals in order of increasing energy.
  • Electrons always fill the lowest available energy level first before moving to higher levels.
  • This allows the atom to achieve maximum stability with the least potential energy.

Energy Levels and Subshells

  • The energy of orbitals depends on both the principal quantum number (n) and the azimuthal quantum number (l).
  • The (n + l) rule is used to determine the order of filling:
    • Orbitals with lower (n + l) are filled first.
    • If two orbitals have the same (n + l) value, the orbital with lower n is filled first.
  • Example: 4s (n + l = 4) fills before 3d (n + l = 5).

Electron-Electron Repulsion and Stability

  • Electrons repel each other due to negative charge.
  • Filling lower energy orbitals first minimizes repulsion because electrons are closer to the nucleus and more tightly bound.
  • This arrangement reduces total energy and enhances the stability of the atom.

Examples of Orbital Filling

  1. Hydrogen (H): 1s¹ → fills the 1s orbital first.
  2. Oxygen (O): 1s² 2s² 2p⁴ → 1s fills first, then 2s, then 2p.
  3. Iron (Fe): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ → 4s fills before 3d due to lower (n + l) energy.

Significance in Chemistry

  1. Electron Configuration: Determines the arrangement of electrons in shells and orbitals.
  2. Chemical Properties: Influences valence electrons and reactivity of elements.
  3. Periodic Table: Explains why elements in the same group have similar properties.
  4. Bonding Behavior: Stable electron arrangements guide ionic and covalent bonding.
  5. Magnetism: The filling order affects the number of unpaired electrons, influencing magnetic properties.

Connection to Other Principles

  • Pauli Exclusion Principle: Limits each orbital to two electrons with opposite spins.
  • Hund’s Rule: Ensures electrons occupy degenerate orbitals singly before pairing.
  • Aufbau Principle: Explains why electrons fill lowest energy orbitals first.

These rules together provide a complete picture of atomic stability and electron arrangement.

Visual Representation

  • Orbitals closer to the nucleus (1s, 2s) fill first.
  • Higher orbitals (3p, 4s, 3d) fill later as electrons are added.
  • This layered filling creates the shell and subshell structure of the atom.
Conclusion

Electrons fill lower energy orbitals first to minimize energy and maximize stability. This principle, based on the Aufbau principle and quantum mechanics, determines electron configurations, chemical reactivity, bonding, and periodic trends. By occupying the lowest available orbitals first, atoms achieve their most stable state, which is fundamental to understanding atomic structure and chemistry.