Rules of Electronic Configuration

Electronic Configuration And Its Rules

As we started digging deeper into the atom, we got to know that electrons can’t exist anywhere around the nucleus. We got to know that electrons are filled up in shells and subshells around the nucleus according to a few rules which are now known as electronic configuration rules.

Before getting into electronic configuration and the rules followed, you need to know about the four quantum numbers used to identify electron, shells, subshells, orbitals, etc. Refer this article to understand all of the mentioned topics.

Electronic Configuration

Electronic COnfiguration

The electronic 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 atomic number of Aluminium is 13 and its electronic configuration is 1s22s22p63s23p1

The electronic configuration of Copper (29) is 1s22s22p63s23p64s23d9

As you could see, there are some rules following which the shells and subshells are filled. So, let us look at the three most important rules for filling electrons in the shells.

Pauli’s Exclusion Principle

The number of electrons that are to be filled in various orbitals is restricted by Pauli’s Exclusion principle. This principle was given by the Austrian scientist Wolfgang Pauli (1926).

According to this principle: No two electrons in an atom can have the same set of four quantum numbers. Pauli exclusion principle can also be stated as: “Only two electrons may exist in the same orbital and these electrons must have opposite spin.” This means that an electron can have the same n, l, and ml but it should at least differ in ms quantum numbers.

This principle imposed on orbitals helps us calculate the capacity of the subshells. For example, the subshell s have only one orbital so it has 2 electrons, similarly, p subshell has 3 orbitals, so it holds 6 electrons.

Hund’s Rule

Hund’s Rule of Maximum Multiplicity deals with the filling of electrons in orbitals of the same shell (The orbitals of the same shell are called degenerate orbitals as they have the same energy).

It states: The pairing of electrons in the orbitals belonging to the same subshell (p, d or f) does not take place until each orbital belonging to that subshell has got one electron each i.e., it is singly occupied.

For example, there are 3 orbitals in p subshell, so, the first orbital will get its second electron only after each of the three subshells have one electron. Basically it means that pairing in degenerate orbitals start only after all of them have 1 electron each

Aufbau’s Principle

This principle deals with the filling of orbitals. The principle states: In the ground state of the atoms, the orbitals are filled in order of their increasing energies.

The orbitals in their increasing order of energy are shown in the below image.

The order of energy of orbitals
The order of energy of orbitals (Source)

So, the electrons are first filled into lower energy orbitals starting from 1s and then gets filled in the energy order. (Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f,5d, 6p, 7s…)

This principle is applicable for most of the elements but exceptions always exist in chemistry :-). Due to stability provided by half-filled or fully filled orbitals and very less energy gaps between a few orbitals like 3d and 4s, there exist some exceptions like Copper and Chromium.

So, those were the three principles for filling electrons in orbitals. I hope you got a clarity regarding this topic. Doubts are highly encouraged in the comments.

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