Electron Configuration

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What are Electron Configurations?

The electron configuration of an element describes how electrons are distributed in its atomic orbitals. Electron configurations of atoms follow a standard notation in which all electron-containing atomic subshells (with the number of electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is 1s22s22p63s1.

However, the standard notation often yields lengthy electron configurations (especially for elements having a relatively large atomic number). In such cases, an abbreviated or condensed notation may be used instead of the standard notation. In the abbreviated notation, the sequence of completely filled subshells that correspond to the electronic configuration of a noble gas is replaced with the symbol of that noble gas in square brackets. Therefore, the abbreviated electron configuration of sodium is [Ne]3s1 (the electron configuration of neon is 1s22s22p6, which can be abbreviated to [He]2s22p6).

Electron Configurations are useful for:

Determining the valency of an element.
Predicting the properties of a group of elements (elements with similar electron configurations tend to exhibit similar properties).
Interpreting atomic spectra.

This notation for the distribution of electrons in the atomic orbitals of atoms came into practice shortly after the Bohr model of the atom was presented by Ernest Rutherford and Niels Bohr in the year 1913.

Writing Electron Configurations

Shells

The maximum number of electrons that can be accommodated in a shell is based on the principal quantum number (n). It is represented by the formula 2n2, where ‘n’ is the shell number. The shells, values of n, and the total number of electrons that can be accommodated are tabulated below.

Shell and ‘n’ value	Max. Electrons in the Electron Configuration
K shell, n=1	2*12 = 2
L shell, n=2	2*22 = 8
M shell, n=3	2*32 = 18
N shell, n=4	2*42 = 32

Subshells

The subshells into which electrons are distributed are based on the azimuthal quantum number (denoted by ‘l’).
This quantum number is dependent on the value of the principal quantum number, n. Therefore, when n has a value of 4, four different subshells are possible.
When n=4. The subshells correspond to l=0, l=1, l=2, and l=3 and are named the s, p, d, and f subshells, respectively.
The maximum number of electrons that can be accommodated by a subshell is given by the formula 2*(2l + 1).
Therefore, the s, p, d, and f subshells can accommodate a maximum of 2, 6, 10, and 14 electrons, respectively.
All the possible subshells for values of n up to 4 are tabulated below.
Principle Quantum Number Value Value of Azimuthal Quantum Number Resulting Subshell in the Electron Configuration
n=1 l=0 1s
n=2 l=0 2s
l=1 2p
n=3 l=0 3s
l=1 3p
l=2 3d
n=4 l=0 4s
l=1 4p
l=2 4d
l=3 4f
Thus, it can be understood that the 1p, 2d, and 3f orbitals do not exist because the value of the azimuthal quantum number is always less than that of the principal quantum number.

Notation

The electron configuration of an atom is written with the help of subshell labels.
These labels contain the shell number (given by the principal quantum number), the subshell name (given by the azimuthal quantum number) and the total number of electrons in the subshell in superscript.
For example, if two electrons are filled in the ‘s’ subshell of the first shell, the resulting notation is ‘1s2’.
With the help of these subshell labels, the electron configuration of magnesium (atomic number 12) can be written as 1s2 2s2 2p6 3s2.

Filling of Atomic Orbitals

Aufbau Principle

This principle is named after the German word ‘Aufbeen’ which means ‘build up’.
The Aufbau principle dictates that electrons will occupy the orbitals having lower energies before occupying higher energy orbitals.
The energy of an orbital is calculated by the sum of the principal and the azimuthal quantum numbers.
According to this principle, electrons are filled in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p…

Pauli Exclusion Principle

The Pauli exclusion principle states that a maximum of two electrons, each having opposite spins, can fit in an orbital.
This principle can also be stated as “no two electrons in the same atom have the same values for all four quantum numbers”.
Therefore, if the principal, azimuthal, and magnetic numbers are the same for two electrons, they must have opposite spins.

Hund’s Rule

This rule describes the order in which electrons are filled in all the orbitals belonging to a subshell.
It states that every orbital in a given subshell is singly occupied by electrons before a second electron is filled in an orbital.
In order to maximize the total spin, the electrons in the orbitals that only contain one electron all have the same spin (or the same values of the spin quantum number).

Representation of electronic Configuration of Atom

The electron configurations of a few elements are provided with illustrations in this subsection.