Aufbau Principle

What is the Aufbau Principle?

According to this, electrons are inserted into atomic orbitals in ascending order of orbital energy. The Aufbau principle states that atomic orbitals that are available and have the lowest energy levels are filled before those that have higher energy levels.

Germanic in origin, the word “aufbau” generally translates to “construct” or “build up.” Below is a schematic showing the sequence in which atomic orbitals are filled. Here, the terms “n” and “l” denote the principal quantum number and azimuthal quantum number, respectively.

The placement of electrons in an atom and the related energy levels can be understood using the Aufbau concept. For instance, the electrical configuration of carbon is 1s22s22p2 and it possesses six electrons.

It’s crucial to remember that each orbital can only accommodate a maximum of two electrons (as per the Pauli exclusion principle). Additionally, the filling of electrons into orbitals within a single subshell must adhere to Hund’s rule, which states that before any two electrons pair up in an orbital, each orbital within a given subshell must be solely occupied by electrons.

Read Also: Henderson-Hasselbalch Equation

Salient Features of the Aufbau Principle

• The Aufbau principle states that electrons first inhabit orbitals with the lowest energies. This suggests that electrons only move into orbitals with higher energy once orbitals with lower energies are fully occupied.
• The (n+l) rule, which states that the orbital’s energy level is defined by the sum of its primary and azimuthal quantum numbers, can be used to identify the sequence in which the energy of orbitals grows.
• Lower orbital energies are correlated with lower (n+l) values. The orbital with the lower n value is said to have lower energy associated with it if two orbitals have equal (n+l) values.
• The order in which the orbitals are filled with electrons is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, and so on.

Exceptions

Instead of [Ar]3d44s2, the electron configuration of chromium is [Ar]3d54s1 (as suggested by the Aufbau principle). The improved stability offered by half-filled subshells and the relatively small energy difference between the 3d and 4s subshells are two reasons for this anomaly, among others.

Below is an illustration of the energy difference between the various subshell.

Half-filled subshells have reduced orbital electron-electron repulsions, which boost stability. Similarly to that, fully filled subshells boost the atom’s stability. As a result, some atoms’ electron configurations defy the Aufbau principle (depending on the energy gap between the orbitals).

Another exception to this rule is copper, which has an electrical configuration matching [Ar]3d104s1. The stability offered by a fully filled 3D subshell can account for this.

Electronic Configuration using the Aufbau Principle

Writing the Electron Configuration of Sulphur

• The atomic number of sulphur is 16, implying that it holds a total of sixteen electrons.
• As per the Aufbau principle, 2 of these electrons are present in the 1s subshell, 8 of them are present in the 2s and 2p subshell, & the remaining are distributed into the 3s and 3p subshells.
• Therefore, the electron configuration of sulphur could be written as 1s²2s²2p⁶3s²3p⁴.

Writing the Electron Configuration of Nitrogen

• Nitrogen has seven electrons (since its atomic number is 7).
• The 1s, 2s, and 2p orbitals are full of electrons.
• Nitrogen’s electron configuration can be expressed as 1s22s22p3.