Electronic Configuration of Elements

The distribution of circular energy levels for s, p, d, and f-block components of the periodic table is known as electronic configuration, general electron configuration, or electronic structure of atoms or ions. In chemistry, the electron is the key to unlocking the chemical universe and learning about the properties and positions of elements in the periodic table. The periodic table elements’ electronic configuration or electron arrangement method is used to better understand chemical properties such as oxidising and reducing properties, oxidation number, ionisation energy, electron affinity, shielding effect, the polarity of chemical bonds, acids and bases properties, and so on. The electronic configuration is connected to periodicity. To put it another way, all chemical and physical qualities are expressions of an element’s electronic configuration.

Distribution of Electrons

• Periodic Table

By classifying the elements, scientists were able to find a systematic approach to organise their information. Not only would it justify known chemical facts regarding elements, but it would also anticipate new ones that could be studied further. The physical and chemical properties of elements are a function of their atomic numbers on a regular basis. The Periodic Law is the result of periodic changes in electronic configurations, which affect the physical and chemical characteristics of elements and their compounds.

The most convenient and frequently used version of the periodic table is the so-called “long form” of the Periodic Table of the Elements. Periods are the lateral rows  (Mendeleev-named series), and groups are the vertical columns. Groups or families of elements with comparable exterior electronic configurations in their atoms are organised in vertical columns.

• Electron Energy Levels

Energy levels are defined as fixed distances in which electrons rotate around the core with a defined amount of energy. With increasing distance from the nucleus, the energy linked to a particular energy level rises. In 1s hydrogen energy levels with electronic configuration 1s1, the hydrogen atom has only one electron.

1. Shells – It is based on the primary quantum number (n), which specifies the highest amount of electrons that can fit inside a shell. The formula 2n2 gives the number of electrons within every shell, where n is the shell number.
2. Subshells – It is built around an azimuthal quantum number, commonly abbreviated as ‘I.’ At the same time, the value of the primary quantum number (n) affects the azimuthal quantum number. It’s these shells where the real electron distribution takes place. The formula 2(2I+1) provides the greatest number of electrons that a subshell can hold. They are of 4 types- s,p,d and f.
3. Magnetic Quantum number – The magnetic quantum number is used to compute the azimuthal component of an orbital’s orientation in space, and it identifies the orbitals present within a subshell.
4. Spin Quantum Number – The rotational motion of an electron is described by the Spin Quantum Number. An electron has both angular and orbital angular momentum as it spins around an axis. The Spin Quantum Number (s) not only has a magnitude (1/2) but also has a direction (+ or -) since angular momentum is a vector.

Principles for Filling the Shells

1. Aufbau Principle: The periodic classification is theoretically supported by the aufbau (build-up) principle and the electrical arrangement of atoms. According to the principle, the electrons inhabit lower energy orbitals first, then higher energy orbits. The term Aufbau comes from the Greek word “Aufbeen,” which means “to build up.”
2. Pauli’s Exclusion Principle: According to the Pauli Exclusion Principle, an orbit with opposite spins can hold a maximum of two electrons. It is based on the idea that even if two electrons have the same azimuthal number, their spins in the same orbit will be opposite.
3. According to the rule, Hund’s Rule: Every orbit is first filled by a free electron before a new electron enters the orbit. This law determines the arrangement of electrons in the subshell.

• Notation of Electronic Configuration

The amount of electrons available in a subshell is represented by notation. The shell number, the name of the subshell, and the overall amount of electrons present in the subshell are all written in superscript. The electrical configuration of oxygen, for example, can be expressed as 1s2 2s2 2p4 .The electronic configuration of Sodium (Na) may be written as 1s2 2s2 2p6 3s1

The line arrangement will be followed by filling the various orbitals by the amount of electrons. The following ordering is used to build electron configuration energy levels: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f.

Conclusion

The number of electrons in each sublevel of an atom’s energy levels in its ground state is known as its electron configuration. When an atom is in its ground state, its electron configuration explains how its electrons fill sublevels. Subshells are partially or fully filled whenever possible because atoms seek the most steady electron configuration. Scientists employ a shorthand notation that begins with the sign of a  neutral gas before the element on the periodic table rather than writing out the entire electron configuration.

In a neutral atom, the atomic number equals the nuclear charge (number of protons) or the number of electrons. The relevance of quantum numbers and electronic arrangements in the periodicity of elements may thus be easily visualised. Noble gases have closed electron shells and very stable electron configurations. Group 17 elements (halogens) have quite high negative electron gain enthalpies because they can take up an electron and achieve stable noble gas electron configurations.