Short Answer

Hund’s rule is a rule in atomic physics that explains how electrons fill orbitals within the same subshell. It states that electrons first occupy different orbitals singly with parallel spins before they start pairing. This helps reduce repulsion and makes the atom more stable.

According to Hund’s rule, electrons always prefer to stay unpaired in different orbitals of the same energy level. Only when each orbital has one electron do they begin to pair up. This rule is very useful in writing correct electron configurations for atoms.

Detailed Explanation :

Hund’s rule

Hund’s rule is an important principle in quantum mechanics that explains the arrangement of electrons in different orbitals belonging to the same subshell. The rule was proposed by German physicist Friedrich Hund. It helps understand why electrons fill orbitals in a particular pattern and how atoms achieve the lowest possible energy state.

Orbitals in the same subshell (such as the three p orbitals or the five d orbitals) have the same energy. These orbitals are called degenerate orbitals. Hund’s rule tells us how electrons distribute themselves among these degenerate orbitals in the most stable way. This principle is essential for writing correct electron configurations, predicting magnetic behaviour, and understanding chemical bonding.

Hund’s rule can be stated simply:

Electrons fill degenerate orbitals singly first, with parallel spins, before pairing begins.

This means that if several orbitals of equal energy are available, electrons will enter them one at a time rather than pairing up immediately. Only after each orbital contains one electron will additional electrons start pairing in those orbitals.

Meaning and explanation of Hund’s rule

Hund’s rule is based on the idea that electrons repel each other because they carry the same negative charge. When electrons fill orbitals singly, they stay as far apart as possible, reducing repulsion and increasing the stability of the atom. When electrons are forced to pair, the repulsion between their opposite spins increases the energy of the system.

Hund’s rule is made up of two key ideas:

1. Maximum multiplicity
   Electrons in degenerate orbitals prefer to have the same spin direction. This arrangement is known as maximum multiplicity. It leads to a lower energy state.
2. Single occupation of orbitals first
   Electrons occupy each orbital of the subshell singly before any orbital gets a second electron.

These ideas ensure that the atom remains in the lowest energy configuration.

Application in different subshells

Hund’s rule applies to any subshell with more than one orbital:

• p subshell → 3 orbitals
• d subshell → 5 orbitals
• f subshell → 7 orbitals

For example, in the p subshell, the three orbitals are px, py, and pz. According to Hund’s rule, if there are three electrons to fill the p subshell, each electron will enter a different orbital with the same spin direction.

Why electrons prefer to remain unpaired

Electrons prefer to remain unpaired in degenerate orbitals because:

• Unpaired electrons with the same spin avoid each other due to repulsion.
• This reduces electron–electron repulsion, lowering the atom’s energy.
• The magnetic fields produced by electrons cancel less, increasing stability.

Pairing electrons too early increases repulsion and makes the atom less stable. Thus, following Hund’s rule ensures the most stable configuration.

Hund’s rule and quantum numbers

Hund’s rule works along with quantum numbers:

• The magnetic quantum number (mₗ) selects which orbital the electron occupies.
• The spin quantum number (mₛ) ensures that electrons first fill orbitals with parallel spins.

By obeying these rules, electrons achieve a stable arrangement.

Importance of Hund’s rule in electron configuration

Hund’s rule plays a major role in determining how electrons fill atomic orbitals:

1. Writing electron configurations correctly
   Without Hund’s rule, many electron configurations would be incorrect, especially in p, d, and f blocks.
2. Predicting magnetic properties
   Elements with unpaired electrons are paramagneticand attracted to magnetic fields.
   Elements with all paired electrons are diamagnetic.
   Hund’s rule helps determine whether an atom has unpaired electrons.
3. Understanding chemical behaviour
   The number of unpaired electrons affects bonding and chemical reactivity. Hund’s rule helps explain why some elements form more bonds than others.
4. Explaining shapes of molecules
   Hybridization and molecular geometry depend on electron configuration, which depends on Hund’s rule.

Examples showing Hund’s rule

Example 1: Nitrogen (N)
Its electron configuration ends with 2p³.
According to Hund’s rule, the three electrons occupy three different p orbitals with parallel spins.

Example 2: Oxygen (O)
Its configuration ends with 2p⁴.
Three of the p orbitals receive one electron each, and the fourth electron pairs in one of the orbitals.

Example 3: Iron (Fe)
d-block elements often have unpaired electrons that follow Hund’s rule.
This helps explain why iron is magnetic.

Hund’s rule and stability of atoms

Hund’s rule also explains some exceptions in electron configurations, especially in transition metals and lanthanides. For example, chromium and copper have electron configurations that maximize stability by increasing the number of unpaired electrons in degenerate orbitals.

Conclusion

Hund’s rule states that electrons fill degenerate orbitals singly with parallel spins before pairing begins. This rule reduces electron repulsion, increases stability, and helps explain the arrangement of electrons in p, d, and f subshells. Hund’s rule is essential for writing electron configurations, understanding magnetism, predicting chemical behaviour, and explaining atomic structure.