A Quick Note on Electronic Configuration

The electronic configuration of an atom or molecule is the distribution of electrons in atoms or molecular orbitals. The orbitals occupied by electrons on an atom are described by the electronic configuration of the atom. This prediction is based on the assumption that electrons are added to an atom one at a time, starting with the lowest energy orbital and progressing to the highest energy orbital until all of the electrons have been placed in approximate orbitals. 

The electronic configuration describes an atom’s orbitals in its ground state, but it may also be used to represent an atom that has ionized into a cation or anion by compensating for electron loss or gain in the subsequent orbitals. Many of an element’s physical and chemical properties can be linked to its electrical arrangement.

What is electronic configuration? 

Electron configurations are a quick way to write down where all of the electrons in an atom are located. As we all know, the positively charged protons in an atom’s nucleus attract negatively charged electrons. Because of their attraction to the protons, these electrons all stick together within the atom, but they also repel each other, leading them to spread out in regular patterns around the nucleus. This produces lovely geometric formations known as orbitals, which represent the different regions around the nucleus that each electron draws out. The Pauli Exclusion Principle, a quantum physics principle that states that no two electrons can ever be in the same position, is the reason why electrons tend to stay in their orbitals rather than stacking on top of one another. The Pauli Exclusion Principle is derived from fundamental physics rules that govern all subatomic particles, rather than only the electrical repulsion of negative electrons.

Each electron in an atom has a unique “address,” which is represented by the orbitals. Consider the electrons to be residents of one of several blocks of studio apartments near a pleasant park. The electrons all want to reside near the park (nucleus), but they can’t. Instead, some electrons are assigned to the apartments closest to the nucleus; but, as the number of electrons wishing to reside near a nucleus grows, some of them will be forced to relocate further away, as apartments closer to the nucleus fill up. This illustrates a periodic table trend: elements with a low atomic number (and consequently fewer electrons) have the majority of their electrons in orbitals close to the nucleus. As we progress down the periodic table, electrons begin to fill orbitals and energy levels further away from the nucleus.

Electronic configuration of carbon

Carbon is a chemical element with the atomic number 6 and the symbol C. It is nonmetallic and tetravalent, which means it has four electrons available for chemical bonding. It belongs to the periodic table’s 14th group. Only roughly 0.025 percent of the Earth’s crust is made up of carbon.

Carbon is denoted by C. Its atomic mass number is 12.01o7u and its atomic number is 6.

Electronic configuration of carbon is [He] 2s2 2p2.

Writing electronic configuration 

Shells

An electron shell surrounds the atomic nucleus in the region of an atom. It’s a group of atomic orbitals with the same value of the primary quantum number, n. One or more electron subshells, also known as sublevels, are included in electron shells.

Shell and n value 

K shell, n = 1

L shell, n =2

M shell, n = 3 

N shell, n = 4

Filling of atomic orbitals 

The Aufbau principle, which is based on Pauli’s exclusion principle, Hund’s rule of maximal multiplicity, and the relative energies of the orbitals, governs the filling of electrons into the orbitals of distinct atoms.

Aufbau principle 

According to this approach, electrons are added one by one to the various orbitals in order of increasing energy, starting with the lowest energy orbital. The energy of various orbitals is arranged in ascending order.

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

Pauli’s exclusion 

“No two electrons in an atom will have the identical value of all four quantum numbers,” according to this concept.

If one electron in an atom has the quantum numbers n = 1, l=0, m = 0, and s = +1/2, no other electron in the atom can have the same four. In other words, two electrons with the same s value cannot be placed in the same 1s orbital.

The orbital graphic does not depict a conceivable electron arrangement.

An orbital can only carry two electrons since there are only two conceivable values of s.

Hund’s rule

This rule governs the filling of electrons in orbitals with the same energy (degenerate orbitals). “Electron pairing in p, d, and f orbitals cannot occur until each orbital of a given subshell possesses one electron or is singly occupied,” according to this law.

This is owing to the fact that identically charged electrons repel each other when they are in the same orbital. This repulsion can be reduced if two electrons inhabit distinct degenerate orbitals and migrate as far apart as possible. In a degenerate collection of orbitals, all unpaired electrons will have the same spin. 

The most important thing to remember is that all of the singly occupied orbitals should have electrons spinning in the same direction, either clockwise or anticlockwise.

Conclusion 

Electron configurations are a quick way to write down where all of the electrons in an atom are located. The electronic configuration describes an atom’s orbitals in its ground state, but it may also represent an atom that has ionized into a cation or anion by compensating for electron loss or gain. Carbon is a chemical element with the atomic number 6 and symbol C. It is nonmetallic and tetravalent, which means it has four electrons available for chemical bonding. Carbon belongs to the periodic table’s 14th group – only 0.025 percent of Earth’s crust is made up of carbon.