General Introduction, Electronic Configuration, Occurrence, and Features of the Phenomenon

General introduction, electronic configuration, occurrence, and features of the phenomenon

 META DESCRIPTION: Electrons orbit the nucleus of an atom, while protons and neutrons are enclosed within it.  An atom’s electronic configuration is a numerical depiction of the distribution of electrons in the orbitals of the atom. This dictates an element’s position in the periodic table and, as a result, its chemical behaviour. It describes how chemical bonds hold atoms together and the strange patterns that may be seen in the periodic table’s rows and columns. 

The arrangement of electrons in energy levels surrounding an atomic nucleus is known as electronic configuration or electronic structure. Electrons occupy multiple levels in the older shell atomic model, from the first shell, K, closest to the nucleus, to the seventh shell, Q, farthest from the nucleus. The K–Q shells are subdivided into a set of orbitals (see orbital) that can each be inhabited by no more than a pair of electrons in a more sophisticated quantum-mechanical model.  

In the shell atomic model, an atom’s electrical configuration can be described by listing the number of electrons in each shell, starting with the first. The 11 electrons in sodium (atomic number 11) are distributed as follows: the K and L shells are totally filled with 2 and 8 electrons, respectively, while the M shell is only half-filled with one electron.

In the quantum-mechanical concept, an atom’s electronic configuration is expressed by listing the occupied orbitals in order of filling, with the number of electrons in each orbital indicated by a superscript.  The electrical configuration of sodium in this notation is 1s2 2s2 2p6 3s1, which is dispersed in the orbitals as 2-8-1. The electrons in excess of the noble gas configuration immediately preceding the atom in the periodic table are frequently listed using a shorthand manner. Because sodium (chemical symbol Ne, atomic number 10) has one more 3s electron than the noble gas neon (chemical symbol Ne, atomic number 10), its shorthand notation is [Ne]3s1.

Electronic configurations are similar among elements in the same periodic table group. Lithium, sodium, potassium, rubidium, caesium, and francium (all alkali metals in Group I) have electrical configurations that display one electron in the outermost (loosely bound) s- orbital. This so-called valence electron is responsible for the alkali elements in Group I having similar chemical properties, such as bright metallic lustre, high reactivity, and good thermal conductivity.

Electronic Configuration

The electron configuration of an atom is the distribution of electrons within the orbits (shells) and subshells of the atom.

The electron configuration of an atom is significant because it aids in the prediction of a substance’s chemical, electrical, and magnetic characteristics.

We can forecast whether two substances will chemically react or not based on the electron configuration of the atom, and if they do, we can also anticipate what kind of reaction will occur and how powerful it will be.

The arrangement of electrons in space around the nucleus is described by the electron configuration of an atom.

The electrons are dispersed over many energy levels. The shell or orbits are the names given to these energy levels. The distribution of electrons over several shells (energy levels) is such that the overall energy of all electrons in an atom remains as low as possible for the atom’s stability. The distribution of electrons at different energy levels is determined by the following rules:

1. ‘2n2‘, where n is an integer and represents the “principal quantum number,” gives the maximum number of electrons in any primary energy level (shell). The value of ‘n’ varies depending on the main energy levels.
2. Subshells are subdivided into each major shell (energy level). Orbitals are the subshells that make up a shell. These subshells/orbitals are denoted by the letters s, p, d, f, etc., together with the orbital quantum number, l = 0, 1, 2, 3, 4,….(n-1) etc. The “primary quantum number” ‘n’ equals the number of subshells in any main shell.
3. The formula 2(2l + 1) governs the maximum electron capacity of subshells.
4. The Construction Principle

“Aufbau” is a German word that means “to construct.” As a result, the “Aufbau Principle” is also known as the “building up principle.” The electrons inhabit the orbitals in a sequence of increasing energy, according to this concept. The following is the increasing energy order of numerous orbitals, as well as the order of occupation —

1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s<5f<6d<7p.

Characteristics of Periods in a Periodic Table

Valence Electrons 

As a period proceeds from left to right, the number of valence electrons in elements increases from 1 to 8, while it increases from 1 to 2 in the first period.

Valency

As you walk from left to right in a period, the valency of elements grows from 1 to 4 and then declines to zero.

Atomic sizes

The atomic size shrinks from left to right over time. As the atomic number increases, the number of protons and electrons increases as well, therefore the extra electrons are added to the same shell. Because of the strong positive charge on the nucleus, electrons are attracted closer to the nucleus, and the atom shrinks in size. As a result, the nucleus has a stronger attraction. As a result, the atomic size decreases.

Metallic Character

As you move from left to right, the metallic character of components reduces while the non-metallic character grows. Sodium, magnesium, and aluminium are among the metals of the third phase. Silicon is classified as a metalloid because its properties are intermediate between those of metals and non-metals. Nonmetallic elements include phosphorus, sulphur, and chlorine.

Group Characteristics in a Periodic Table

Valence Electrons

A periodic table group’s elements all have the same number of valence electrons. For example, lithium, sodium, and potassium all have one valence electron in their atoms and are members of group 1 of the periodic table. Lithium, sodium, and potassium atoms can all easily lose their one valence electron, resulting in Li, Na, and K ions with one unit positive charge, respectively.

Valency

Because the number of valence electrons that define valency is the same for all elements in a group, they all have the same valency. For example, lithium, sodium, and potassium all have one valence electron, hence all elements in group 1 have the same valency.

Atomic sizes

As one proceeds down the periodic table, the size of atoms, or atomic size, increases. The size of the atoms gradually rises from lithium to francium as we move down in group 1 from top to bottom. A new shell of electrons is added to the atoms every time we move from the top to the bottom of a group.

Metallic Character

As you progress from top to bottom, the metallic character of elements grows while the non-metallic character declines. The elements in the group’s bottom half have the most metallic appearance. From lithium to francium, for example, the metallic character of group 1 increases. As we advance down the periodic table, one more electron shell is added, and the size of the atoms increases. Because the number of valence electrons that define valency is the same for all elements in a group, they all have the same valency.

CONCLUSION

We looked at why alkali and alkaline earth metals are called s-block elements. The abnormal behaviour of the group’s initial element was investigated, as well as the diagonal link between second and third-period members. We now understand the physical and chemical properties of alkali and alkaline earth metals, as well as their applications. 

The arrangement of electrons in energy levels surrounding an atomic nucleus is known as electronic configuration or electronic structure.

It also describes how chemical bonds hold atoms together and the strange patterns that may be seen in the periodic table’s rows and columns.