Short Answer

Electron configuration is the arrangement of electrons in different shells, subshells, and orbitals of an atom. It shows how electrons are distributed around the nucleus according to their energy levels. This arrangement follows specific rules of quantum mechanics.

Electron configuration helps explain the chemical behaviour, bonding nature, and physical properties of elements. It is written using a notation such as 1s² 2s² 2p⁶, which tells us how many electrons are present in each orbital or subshell.

Detailed Explanation :

Electron configuration

Electron configuration refers to the systematic arrangement of electrons in an atom according to increasing energy levels. It tells us how electrons fill orbitals, subshells, and shells. This arrangement is governed by quantum numbers, which describe the energy, shape, and orientation of orbitals. Electron configuration is one of the most important concepts in atomic structure because it determines nearly all physical and chemical properties of elements, such as reactivity, valency, magnetism, and bonding behaviour.

Electrons do not revolve in fixed circular paths around the nucleus, as suggested by older models. Instead, according to quantum mechanics, they occupy orbitals, which are defined as three-dimensional regions where the probability of finding an electron is high. These orbitals are grouped into subshells (s, p, d, f) and shells (n = 1, 2, 3, …). The way electrons fill these orbitals is described by electron configuration.

Quantum basis of electron configuration

Electron configuration is based on four quantum numbers:

• Principal quantum number (n)
• Azimuthal quantum number (l)
• Magnetic quantum number (mₗ)
• Spin quantum number (mₛ)

The values of these quantum numbers determine which orbitals exist and how many electrons each orbital can hold. The maximum number of electrons in an orbital is two, and they must have opposite spins due to the Pauli Exclusion Principle.

Rules governing electron configuration

Three main rules guide the filling of electrons in orbitals:

1. Aufbau principle
   Electrons fill orbitals in order of increasing energy. Lower-energy orbitals fill first, then higher ones.
   Example: 1s fills before 2s, 2p fills before 3s.

The general filling order is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s …

2. Pauli Exclusion Principle
   No two electrons in the same atom can have identical sets of four quantum numbers.
   Therefore, each orbital can hold only two electrons, and they must have opposite spins.
3. Hund’s rule
   Electrons fill degenerate orbitals singly first, with parallel spins, before pairing begins.
   For example, the three p orbitals get one electron each before any of them receives a second electron.

These three rules ensure that atoms achieve the most stable electron arrangement possible.

Structure of shells and subshells

Electron configuration depends on the structure of shells and subshells:

• s subshell → 1 orbital → 2 electrons
• p subshell → 3 orbitals → 6 electrons
• d subshell → 5 orbitals → 10 electrons
• f subshell → 7 orbitals → 14 electrons

Thus, the number of electrons each subshell can hold is fixed.

Writing electron configurations

Electron configuration is written using a specific notation:

• The number (1, 2, 3, …) represents the shell (principal quantum number n).
• The letter (s, p, d, f) represents the subshell.
• The superscript number shows how many electrons are in that subshell.

Example:
1s² 2s² 2p⁶ 3s² 3p⁶

This configuration tells us that:

• 2 electrons are in the 1s subshell
• 2 electrons are in 2s
• 6 electrons are in 2p
• 2 electrons are in 3s
• 6 electrons are in 3p

This is the electron configuration of argon.

Electron configuration and the periodic table

The periodic table is built around electron configurations. Elements are arranged by increasing atomic number, and their positions tell us which subshell is being filled.

• s-block elements fill s orbitals
• p-block elements fill p orbitals
• d-block elements (transition metals) fill d orbitals
• f-block elements fill f orbitals

This arrangement explains why elements in the same group have similar properties—they have similar valence electron configurations.

Importance of electron configuration

Electron configuration plays a key role in:

1. Chemical bonding
   Valence electrons determine whether an element forms ionic, covalent, or metallic bonds.
2. Reactivity of elements
   Elements with nearly filled or nearly empty outer shells are highly reactive.
3. Magnetic properties
   Unpaired electrons cause paramagnetism, while paired electrons cause diamagnetism.
4. Ion formation
   Atoms gain or lose electrons to achieve stable configurations like noble gases.
5. Spectral lines
   Electron transitions between orbitals produce characteristic emission or absorption lines.

Exceptional configurations

Some elements, especially transition metals, show exceptions because half-filled and fully filled subshells are more stable.
Examples:

• Chromium: 3d⁵ 4s¹ (instead of 3d⁴ 4s²)
• Copper: 3d¹⁰ 4s¹ (instead of 3d⁹ 4s²)

These exceptions still follow the principle of minimum energy.

Conclusion

Electron configuration describes how electrons are arranged in an atom’s orbitals, subshells, and shells. It is based on quantum numbers and guided by the Aufbau principle, Pauli Exclusion Principle, and Hund’s rule. Electron configuration helps explain atomic structure, periodic trends, bonding behaviour, and chemical reactivity. It is one of the foundational concepts of modern physics and chemistry.