Electronic Configuration

Electronic configuration means the arrangement of electrons in different energy levels (shells and subshells) of an atom. It tells us how electrons are distributed around the nucleus. This arrangement is important because the chemical properties and reactivity of an element depend on how its electrons are placed. Understanding electronic configuration helps us explain valency, bonding, and periodic trends in the periodic table.

The representation of electrons distributed in the atomic shells of an element is known as the electronic configuration. The electrons are mathematically located in these subshells, and the notations help in locating the position of these electrons as well as the electronic configuration. 

Example: the specific notation for germanium (Ge) would be,1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p².

Distribution of Electrons in Atomic Orbits

The distribution of electrons in atomic orbitals explains how electrons are arranged within an atom. Electrons do not move randomly around the nucleus; they occupy specific regions of space called orbitals. Each orbital has a fixed energy and can hold a limited number of electrons. The arrangement of electrons in these orbitals follows certain rules, which help us understand the structure and chemical behaviour of elements.

1) Shells

Shells are the different energy levels in an atom where electrons are arranged around the nucleus. These energy levels represent the distance of electrons from the nucleus. Electrons present in different shells have different energies.

• The shell nearest to the nucleus has the lowest energy.
• As the shell number increases, the energy of electrons increases.
• Electrons first fill the inner shells before moving to outer shells.
• The maximum number of electrons that can be present in a shell is given by the formula:

N = 2n^2

where:

n = shell number

Example:

2) Subshells

Each shell is further divided into smaller parts called subshells. These subshells help in arranging electrons more accurately within a shell. They are represented by the letters s, p, d, and f. Each subshell has a specific shape, energy, and capacity to hold electrons.

a) Number of Orbitals in Each Subshell

Each subshell contains a certain number of orbitals:

• s subshell → 1 orbital
• p subshell → 3 orbitals
• d subshell → 5 orbitals
• f subshell → 7 orbitals

b) Maximum Number of Electrons in Each Subshell

Each orbital can hold a maximum of 2 electrons.

So:

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

c) Azimuthal Quantum Number (l)

• The azimuthal quantum number explains the shape of the orbital and identifies the subshell in which an electron is present.
• It is also called the angular momentum quantum number.
• It determines the subshell (s, p, d, f).
• It tells us the shape of the orbital.
• It depends on the principal quantum number n.

Example:

In the second shell (n = 2) I= 0 ,I=1:

It has two subshells:

• 2s
• 2p

2s can hold 2 electrons.
2p can hold 6 electrons.

So total electrons in second shell = 8.

3) Notation

Notation is the representation of the number of electrons present in the subshell. It is written with the shell number, name of the subshell, and a total number of electrons present in the subshell.

Example: the electronic configuration of Oxygen can be written as 1s² 2s² 2p⁴.

The below table consists of the electronic configuration of the first 20 elements.

Filling of Atomic Orbitals

Electrons do not fill orbitals randomly; they occupy them in a specific order based on their energy levels. This order of filling is explained by certain rules and principles, which help us write the correct electronic configuration of an atom.

1) Aufbau Principle

The principle states that the electrons will occupy the orbits with lower energy and then the orbits with higher energy levels. The word Aufbau is a German word which means “building up.”

Electrons fill orbitals in order of increasing energy.

Order of Filling of Orbitals

The order is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

4s is filled before 3d because 4s has slightly lower energy than 3d.

The energy of an orbital depends on:

• Principal quantum number (n)
• Azimuthal quantum number (l)

The energy order is decided by the (n + l) rule:

• The orbital with smaller (n + l) value fills first.
• If (n + l) is same, the orbital with smaller n fills first.

Example:

Oxygen (Atomic Number = 8)

Oxygen has 8 electrons.

1s → can hold 2 electrons
2s → can hold 2 electrons
2p → can hold 6 electrons

Step-by-step filling:

1s²→ 2 electrons
2s²→ 2 electrons
2p² → 4 electrons

So electronic configuration is:

1s²  2s²  2p⁴

2) Pauli Exclusion Principle

The Aufbau Principle explains the order in which electrons fill different orbitals based on energy. However, it does not tell us how many electrons can stay in one orbital or how electrons are arranged within orbitals of the same energy this was explained by this principle.

An orbital can hold a maximum of 2 electrons, and they must have opposite spins.

If two electrons are present in the same orbital:

• One must have spin up (↑)
• The other must have spin down (↓)

Example:

Consider the 1s orbital.

It can hold only 2 electrons:

1s²

Orbital diagram:

↑↓

You cannot put a third electron in the same orbital

3) Hund's Rule

Pauli Exclusion Principle could not explain how are electrons arranged in orbitals of the same subshell. This is explained by Hund’s Rule. This states that electrons first occupy empty orbitals singly before pairing up in any orbital. It's the rule that determines the order of electrons in the sub-shell.

First fill each orbital with one electron, then start pairing.

Example: p Subshell (3 Orbitals)

A p subshell has 3 orbitals.

If we have 3 electrons:

Correct filling (Hund’s Rule):

↑ ↑ ↑

Wrong filling :

↑↓ ↑ —

Electrons should not pair until all orbitals have one electron.

Electronic Configuration of Elements

Electronic configuration represents the way in which electrons are filled inside the orbital of any atom. Electronic configuration of various elements are:

1) Carbon

Carbon has six electrons.

Electronic configuration: 1s ²  2s ²  2p ²

2) Neon

Neon has ten electrons.

Electronic configuration: 1s²  2s²  2p⁶

3) Magnesium

Magnesium has twelve electrons.

Electronic configuration: 1s²  2s²  2p⁶  3s²

Related Articles:

• de-Broglie
• Heisenberg Uncertainty Principle