The electron configuration is a representation of how electrons are arranged in an atom’s energy levels or orbitals. It follows certain rules based on the principles of quantum mechanics and the Pauli exclusion principle, which states that no two electrons in an atom can have the same set of quantum numbers. However, there are two exceptions to the rules of electron configurations known as Hund’s rule and the Aufbau principle.
- Hund’s Rule: According to Hund’s rule, when electrons are filling up degenerate orbitals (orbitals with the same energy level), they prefer to occupy different orbitals with the same spin before pairing up. This means that electrons will occupy empty orbitals in the same subshell with the same spin before they start pairing up in opposite-spin orbitals. This results in unpaired electrons in degenerate orbitals, which gives rise to magnetic properties in atoms and ions.
For example, let’s consider the electron configuration of Chromium (Cr) with atomic number 24. Based on the Aufbau principle, the electron configuration of a chromium atom would be: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵. However, according to Hund’s rule, the five 3d orbitals would each have one electron with the same spin (upward arrow) before any of them pair up. This results in a half-filled 3d subshell with five unpaired electrons, which is a more stable arrangement than having four paired electrons and one unpaired electron in the 3d subshell.
- Aufbau Principle: The Aufbau principle states that electrons fill up the lowest energy orbitals first before moving to higher energy orbitals. This means that electrons will fill up orbitals in order of increasing energy levels or increasing values of n+l, where n is the principal quantum number and l is the azimuthal quantum number. However, there are exceptions to this principle when it comes to the relative energies of orbitals.
For example, let’s consider the electron configuration of Copper (Cu) with atomic number 29. Based on the Aufbau principle, the electron configuration of a copper atom would be: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰ 4p⁶. However, experimental observations have shown that the 3d subshell has slightly lower energy than the 4s subshell in copper. As a result, one electron from the 4s subshell is promoted to the 3d subshell, leading to the electron configuration of Cu as: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s⁰. This is known as the “exceptional configuration” or “anomalous electron configuration” of Copper.
Similarly, another example is the electron configuration of Chromium (Cr) as mentioned earlier. Based on the Aufbau principle, the electron configuration of chromium would be expected to be 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴. However, due to the stability associated with half-filled or fully-filled subshells, one electron from the 4s subshell is promoted to the 3d subshell, resulting in the observed electron configuration of Cr as: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵.