Electronic configurations

The distribution of electrons into an atom’s orbitals is known as its electronic configuration. It is the atomic number, which is equal to the number of protons in an atom, that determines how many electrons an atom has. Nuclear nuclei may or may not be the same number of protons and electrons in an atom. An atom’s nucleus consists of protons and neutrons, which have little or no role in regulating chemical reactions. When it comes to the atom’s chemical reactions, electrons are the ones that are outside the nucleus, and their perfect distribution inside an atom is critical. To comprehend the chemical reactions that an atom or its associated molecules participate in, it is necessary to know the precise distribution of electrons in an atom.

Electronic configuration of Elements

An element’s atomic orbitals are described by its electron configuration of elements. There is a standard nomenclature for the arrangement of electron-containing subshells of atoms in which the number of electrons they contain is written in superscript (the number of electrons each subshell contains). Standard notation, however, tends to result in long electron configurations (especially for elements having a relatively large atomic number). In some instances, a shortened or condensed notation may be used instead of the regular notation. Noble gas symbols are used instead of subshells in shortened notation when a noble gas’s electrical configuration can be summarised in a single symbol.

Uses of the electronic configuration of elements 

Electron Configurations are useful for:

• Method for ascertaining an element’s valency.
• When a bunch of elements’ qualities can be predicted (elements with similar electron configurations tend to exhibit similar properties).
• Interpreting the atomic spectra.

Ground state electronic configuration

The lowest-energy and most stable structure is called ground state electronic configuration. A higher energy configuration is an excited state configuration (it requires energy input to create an excited state). Understanding molecule bonding, characteristics, and structures begin with an understanding of ground-state electron configurations. The ground-state electron configuration explains a wide range of atomic properties, from the number of electrons in an atom to its orbitals and hybridisation. The periodic table can only be fully grasped when one has a firm grasp of electron configuration.

Why are electronic configurations important?

To better understand the chemical properties of elements, electron configurations can be used to determine the valence electrons of the atoms. In addition, it aids in the division of elements into several groups (such as the s-block elements, the p-block elements, the d-block elements, and the f-block elements). As a result, it’s easier to investigate the qualities of various elements as a whole.

The electronic configurations of the noble gases are listed below.

• Helium (He) – 1s2
• Neon (Ne) – [He]2s²2p6
• Argon (Ar) – [Ne]3s²3p6
• Krypton (Kr) – [Ar]3d104s²4p6
• Xenon (Xe) – [kr]4d¹05s25p6
• Radon (Rn) – [Xe]4f135d106s26p6

Writing Electron Configurations of all elements 

Here are a few things to keep in mind to write Electronic Configuration of all elements.

Shells

It is based on the fundamental quantum number that determines how many electrons may fit in a shell (n). 

Subshells

• The azimuthal quantum number (‘l’) is used to determine the number of electrons in each subshell.
• On the other hand, this quantum number depends on the value of n. As a result, there are four distinct subshells that can be created when n is equal to 4.
• At this point, n is equal to four. There are four distinct subshells, each of which is designated after a l=0, l=1 to l=3, and l=2 to l=3.
• The formula 2*(2l + 1) gives the maximum number of electrons that can fit in a subshell.
• Consequently, each subshell can hold up to 2, 6, 10, or 14 electrons, based on the electron density.

Filling of Atomic Orbitals

Aufbau Principle

• “Aufbeen,” which means “to build up,” is the German word for this concept.
• This theory states that electrons will first occupy orbitals with lower energies and then move on to orbitals with higher energies.
• Principal and azimuthal quantum numbers are used to compute an orbital’s energy.

Pauli Exclusion Principle

• The Pauli exclusion principle asserts that only two electrons with opposite spins can fit in an orbital at once.
• No two electrons in the same atom have the identical values of all four quantum numbers, according to this rule.
• This means that two electrons must have opposing spin if their primary, azimuthal, and magnetic numbers are the same one for the other.

Hund’s Rule

• A subshell’s orbitals are filled sequentially by this rule.
• There must be a single electron in every orbital in a particular subshell before a subsequent electron is added to an orbital.
• When only one electron is present in an orbital, all of the electrons have the same spin (or the same values of the spin quantum number).

Conclusion

Higher-level chemistry courses require an understanding of electronic configuration. For a complete understanding of the subject, this is the most critical section in the entire chemistry curriculum. When an atom has the correct electrical configuration, it can engage in many chemical reactions, as well as unite with other atoms to produce more complex substances. Electron configurations are a convenient way to record the exact locations of all of an atom’s electrons, as we learn in this book’s chapter. Despite their mutual attraction to the protons, these electrons spread out in predictable patterns around the nucleus because of their mutual attraction to one other.