Electronic Configuration

The electronic configuration is the orbital description of the electrons in an unexcited atom. That is the arrangement of electrons in various atomic orbitals in an atom. Electronic configuration of atoms is based on three important rules which you studied in first year. They are,

i. Pauli's Exclusion Principle

ii. Afbau Principle

iii. Hund's Rule of Maximum Spin Multiplicity

i. Pauli's Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers.

ii. Afbau Principle

Electrons are arranged in various atomic orbitals in the increasing order of their energy.

iii. Hund's Rule of Maximum Spin Multiplicity

Pairing of electrons in degenerate orbitals is not possible until all available orbitals contain one electron each.

According to Pauli's exclusion principle, an orbital can contain a maximum of two electrons and these two electrons must have opposite spin.

The Aufbau principle states that in the ground state of an atom, the orbital with a lower energy is filled up first before the filling of orbitals with higher energy commences.

The increasing order to energy of various atomic orbitals is;

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p etc.

This can be schematically represented as,

Hund's rule deals with the filling of electrons in orbitals having same energy levels. According to this rule electron pairing in degenerate orbitals will not take place until all the available orbitals of a given subshell contain one electron each. For example consider the electron filling in B, C, N and O.

According to Hund's rule, each of the three p-orbitals must get one electron each of parallel spin before each of them gets the second electron of opposite spin.

We can predict how atoms will react on the basis of electronic configuration. Here the importance goes to outer electrons rather than inner electrons. So the inner electron shell notation can be given by replacing the long hand orbital description with the symbol of a noble gas in brackets. This method simplifies the description for bigger atoms. For example the electronic configuration of Be is 1s² 2s² but we write [He]2s².

Here Cr and Cu (Z = 24 and Z = 29) do not follow the general trend.

Z = 24 :      Cr -         1s² 2s² 3s² 3p⁶ 4s² 3d⁴
                                  (Expected configuration)

                                  1s² 2s² 3s² 3p⁶ 4s¹ 3d⁵
                                   (Actual configuration)

Z = 29         Cu -  1    1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹
                                    (Expected configuration)

                                   1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
                                     (Actual configuration)

These abnormalities are due to the extra stability of exactly half-filled and completely filled orbitals ( d⁵, d¹⁰, f⁷, f¹⁴ etc.)

Stability of Completely Filled and Half Filled Subshells

Half filled configurations are more stable due to their greater electron exchange energy. The electrons with parallel spins present in the degenerate orbitals can exchange their positions. The energy released during this exchange is known as exchange energy. In d⁵ configuration there are maximum number of electron exchanges as shown bellow: