Electronic Configuration of Elements

Atomic orbitals define the distribution of electrons in an element’s electron configuration. The electronic configuration of atoms follows a standard nomenclature in which all electron-containing atomic subshells appear in a sequence.

The electronic configuration of elements in different orbitals can be represented as nlx, where:

• n = number of principal shells
• l = symbol of the orbital
• x = number of electrons in the particular orbital

For example, the electron configuration of oxygen (‘O’; Atomic Number = 8) is 1s22s22p4.

Electronic configurations can be useful in: 

• Determining the valency of an element
• Predicting the characteristics of a set of elements with similar electronic configurations
• The study and interpretation of atomic spectra

This nomenclature for the distribution of electrons in atomic orbitals became prevalent after Ernest Rutherford and Niels Bohr published their Bohr model of the atom in 1913.

How to do Electronic Configuration?

The electronic configuration describes where electrons are most likely to be found in an atom. The periodic table’s element blocks help determine electron orbitals. As an example,

• S-block contains Group 1 [alkali metals] and Group 2 [alkaline earth metals]
• D-block contains Groups 2 to 12 metals
• P-block includes Group 13 to Group 18 metals
• F-block contains the lanthanides and actinides

Electron Shells

The Bohr model of atom contributed to the conception of electron configuration. Despite the advances in the quantum-mechanical understanding of electrons, the concept of shells and subshells are still widely used.

An electron shell is an orbit in which electrons revolve around the nucleus of an atom. The highest number of electrons that may be accommodated in a shell is determined by the principal quantum number (n). The formula for the shell number is 2n2, where ‘n’ represents the number of shells. As a result, the total number of electrons that can be fit into different shells are:

• K shell (n=1); Maximum no. of electrons = 2*12 = 2
• L shell (n= 2); Maximum no. of electrons = 2* 22 = 8
• M shell (n=3); Maximum no. of electrons = 2* 32 = 18
• N shell (n=4); Maximum no. of electrons = 2*42 = 32

Subshells 

Every shell is consists of at least one subshell, which itself is made up of atomic orbitals. The formula 2*(2l + 1) gives the maximum number of electrons that may be accommodated by a subshell. 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.

• s orbital – can accommodate a maximum of 2 electrons
• p orbital – can hold 6 electrons
• d orbital – can hold 10 electrons
• f orbital – can hold 14 electrons

Based on different rules and principles, the guidelines for determining the electrical configuration of elements may be summarized as follows:

1. The Aufbau Principle: It states that the orbitals are filled in increasing order of their energies: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, and so on.
2. Hund’s Rule: According to Hund’s Rule, before any orbital in a subshell becomes doubly occupied, it must be singly occupied with one electron.
3. The Pauli Exclusion Principle: According to this principle, no more than two electrons in the same atom can have the same quantum number for all four. In other words, no more than two electrons may occupy the same orbital simultaneously, and they must spin in opposite directions.

Some Practice Examples

Alternatively, you can write the noble gas symbol for an element, such as Lead (Pb) and just add the extra information, i.e., [Xe] 6s2 4f14 5d106p2.

Electronic configuration of chlorine (Cl)

Chlorine is the second-lightest halogen after fluorine. Its atomic number is 17. Therefore, its electronic configuration will be: 1s2, 2s2, 2p6, 3s2, 3p5

The electronic configuration of noble gases are:

• Helium (He): 1s2
• Neon (Ne): [He] 2s2 2p6
• Argon (Ar): [Ne] 3s2 3p6
• Krypton (Kr): [Ar] 3d10 4s2 4p6
• Xenon (Xe): [Kr] 4d10 5s2 5p6
• Radon (Rn): [Xe] 4f14 5d10 6s2 6p6

The Exception Rule

When the electron orbitals are filled or half-filled, then the electronic configurations of elements are in their most stable form. When the energies of the two subshells differ in some elements, an electron may transfer from one to the other, resulting in a symmetrical distribution of electrons in distinct orbitals of the higher energy subshell.

Symmetrical Distribution

As the name implies, symmetrical distributions are electron distributions in which the orbitals of the same subshell are either fully or partly filled, and as such, have a more stable structure. The symmetry results from a balanced distribution of electrons among orbitals.

There is an exception in the d block, notably the groups containing Chromium and Copper, in terms of how they are filled. Due to the narrow energy difference between the 3d and 4s orbitals, this structure violates the Aufbau principle. The fully filled d-orbital arrangement is more stable than the partially filled configuration.

Examples

• Electronic configuration of copper (Cu)

The expected electronic configuration of copper is: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9. However, the actual configuration of Copper is: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10

• Electronic configuration of chromium (Cr)

The expected electronic configuration of chromium is 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4. The actual electronic configuration of chromium is 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5

Periodic Properties

The periodic table displays the elements in increasing order of atomic number. We can also predict the physical, chemical, and atomic properties of these elements by using the periodic law and table formation.

• The electronegativity of the elements increases from left to right and bottom to top in the periodic table
• It is the ability of an atom to attract electrons to itself
• The electron affinity grows from left to right and bottom to top of the periodic table, much like electronegativity
• The total energy received or released by an element because of the addition of one electron is known as electron affinity
• The size of atoms (atomic size) grows from top to bottom in the periodic table, as understood from the orbital diagrams

Conclusion

These electrical configurations aid in the grouping of elements into distinct blocks. It also determines the valence electrons of atoms, which offers details about the chemical behaviour of elements. Understanding the basics of the element’s configuration is important for the overall understanding of the structure of atoms and their nature.