In the realm of atomic and molecular physics, there exist certain principles that govern the behaviour and arrangement of electrons within atoms and molecules. One such fundamental principle is known as Hund’s Rule, named after the German physicist Friedrich Hund, who proposed it in 1927. Hund’s Rule provides insight into the distribution of electrons in atomic orbitals, shedding light on the stability and energy states of atoms and molecules.
Statement of Hund's Rule:
Hund’s Rule can be stated as follows:
“Electrons tend to occupy separate orbitals within a subshell of an atom before they pair up in the same orbital, and they do so to maximise the total spin quantum number (S) and consequently, the overall stability of the atom or molecule.”
In simpler terms, when filling orbitals within a subshell, electrons prefer to occupy separate orbitals with parallel spins (i.e., spins in the same direction) rather than pairing up in the same orbital until necessary. This results in the maximum possible value for the total spin quantum number, which enhances the stability of the atom or molecule.
Explanation of Hund's Rule:
To understand Hund’s Rule better, let’s consider the electron configuration of carbon (C) as an example. Carbon has six electrons, distributed in the 1s²2s²2p² configuration. According to Hund’s Rule, when filling the 2p subshell, the electrons will occupy separate orbitals with parallel spins (↑↑↑) before pairing up in the same orbital. This configuration (↑↑↓↑↓) minimises electron-electron repulsions, leading to lower energy and increased stability.
Hund’s Rule also plays a crucial role in determining the magnetic properties of atoms and molecules. The maximum spin state resulting from electrons occupying separate orbitals with parallel spins contributes to the overall magnetic moment.
Examples in Atomic Physics:
- Helium Atom (He): Helium has two electrons, and its ground state configuration is 1s². Both electrons occupy the 1s orbital, each with opposite spins (↑↓), as dictated by the Pauli Exclusion Principle. Here, Hund’s Rule doesn’t apply since there is only one orbital available in the 1s subshell.
- Oxygen Atom (O): Oxygen has eight electrons, and its ground state configuration is 1s²2s²2p⁴. When filling the 2p subshell, each electron occupies a separate orbital with parallel spins (↑↑↑↑) following Hund’s Rule, before pairing up in the same orbitals (↑↑↓↓).
Examples in Molecular Physics:
- Nitrogen Molecule (N₂): Nitrogen molecules consist of two nitrogen atoms sharing a triple bond. Each nitrogen atom has seven electrons, and the molecular orbital diagram for N₂ is built from the atomic orbitals of two nitrogen atoms. The molecular orbitals are filled according to Hund’s Rule, with electrons occupying separate orbitals before pairing up, resulting in a stable configuration.
- Oxygen Molecule (O₂): Similar to nitrogen, oxygen molecules (O₂) contain two oxygen atoms sharing a double bond. Each oxygen atom has eight electrons. When forming O₂, the molecular orbitals are filled in accordance with Hund’s Rule, ensuring maximum stability and minimising electron-electron repulsions.
Conclusion :
In conclusion, Hund’s Rule is a fundamental principle in atomic and molecular physics that governs the arrangement of electrons within atoms and molecules. By obeying this rule, electrons maximise the total spin quantum number and contribute to the stability and magnetic properties of the system. Understanding Hund’s Rule is essential for comprehending the electronic structure and behaviour of atoms and molecules, laying the groundwork for various applications in chemistry, materials science, and beyond.
