What is Electronic Configuration?
In atomic physics and quantum chemistry, electron placement is the distribution of electrons in an atom or molecule (or other physical structure) in the orbit of the atom or molecule. For example, the electron configuration of a neon atom is 1s2 2s2 2p6. This means that the 1s, 2s, and 2p subshells are occupied by 2, 2, and 6 electrons, respectively. Moving diagonally through the periodic table, the elements show certain similarities, but these are far less pronounced than the similarities within the group. The diagonal relationship is especially noticeable in the elements of the 2nd and 3rd periods of the periodic table.
The electron configuration of the element represents how the electrons are distributed in their atomic orbitals. The electron configuration of an atom follows the standard notation in which an atomic subshell containing all electrons (including the number of electrons in the superscript) is arranged in sequence. For example, the electron configuration of sodium is 1s2 2s2 2p6 3s1. We shall find out the rules for filling electrons in orbitals ahead.
Electronic Configuration of Elements
Shell
The maximum number of electrons that can be accommodated in a shell depends on the principal quantum number (n). This is expressed by equation 2n2. Here, "n" is the shell number. The shell, the value of n, and the total number of electrons that can be accommodated are shown below.
The table below shows the maximum number of electrons that a shell can have, and shell and 'n' value shell.
Number of Electrons in a Shell
Subshell
The subshell to which the electrons are distributed is based on the orbital angular momentum (indicated by "l").
This quantum number depends on the value of the principal quantum number n. Therefore, if the value of n is 4, there are 4 different subshells possible. When n = 4, the subshells correspond to l = 0, l = 1, l = 2, and l = 3, and are called s, p, d, and f subshells, respectively. The maximum number of electrons that a subshell can contain is given by Equation 2 * (2l + 1).
Azimuthal Quantum Numbers
Therefore, the s, p, d, and f subshells can hold up to 2, 6, 10, and 14 electrons, respectively. Above are all possible subshells with values from n to 4; main quantum number value orbital angular momentum value subshell resulting from electron configuration.
Therefore, it is understood that the 1p, 2d, and 3f orbitals do not exist because the orbital angular momentum value is always smaller than the main quantum number.
Notation
The electron configuration of atoms is described using subshell labels. These labels contain the shell number (indicated by the principal quantum number), the subshell name (indicated by the orbital angular momentum), and the total number of electrons in the superscripted subshell.
For example, if two electrons are filled in the "s" subshell of the first shell, the resulting notation would be "1s2".
Aufbau Principle Example
To generate the ground-state electron configuration of any element, we need to know how the atomic sublevels are organised in order of increasing energy.
Electrons are added to atomic orbitals in order from low energy to high energy according to the Aufbau principle.
The lowest energy sublevel is always the 1s sublevel and consists of one orbit.
The only electron in a hydrogen atom occupies a 1s orbital when the atom is in the ground state.
When you move to an atom with multiple electrons, those electrons are added to the next lower sublevel (2s, 2p, 3s, etc.).
After the 3p sublevel, it seems natural that the 3d sublevel has the next lowest energy.
However, the 4s sublevel has slightly lower energy than the 3d sublevel. So it fills up first. This is due to the “n+l” value rule, which states that higher the value of n+l, higher will be the energy.
After filling the 3d sublevel, 4p, 5s, 4d follow.
Note that the 4f sublevel is only filled immediately after the 6s sublevel.
Now the next question is “How to fill electrons in orbitals”
Rules for Filling of Electrons in Orbitals
The three rules for filling the orbit are:
Rule 1:
The lowest energy orbit must be filled first. The filling pattern is as follows:
1s, 2s, 2p, 3s, 3p, 4s, 3d and so on.
The orbit of the subshell is degenerate.
That is, the entire subshell must be filled before filling the next orbit.
This depends on the energy of the lower shell.
Rule 2:
According to the Pauli exclusion principle, the number of electrons allowed per orbit is 2, and they must have opposite spins.
If one electron in the orbit is spinning clockwise, the other electron in the orbit must be spinning counterclockwise.
Such electrons are called unpaired electrons.
Rule 3:
According to Hund's rule, the subshell provides maximum stability when the number of unpaired electrons is maximum, that is, when the spin directions are the same.
This is due to the degeneracy of the orbit.
Therefore, all orbits in the same subshell have the same energy. A pair of electrons occurs when all orbitals are filled individually.
These are also known as the order of filling orbitals.
Summary
The structural principles establish some basic criteria for filling the orbits of atoms. Ground state atoms have the lowest energy and are the most stable. The structural principle regulates the filling of orbitals in the ground state of an atom. These principles are based on Pauli exclusion principle, Hund's rule of maximum multiplicity, and orbital relative energy. According to the principle, electrons are introduced into different orbits as the order of energy increases. The electrons first enter the lowest available energy orbit and then move to the higher energy orbit when the low energy orbit is filled.
FAQs on Rules for Filling of Electron in Orbitals for JEE
1. Is the Aufbau Principle true and which elements are exceptional of the Aufbau Principle?
The Aufbau Principle is used to predict the electron configuration of an atom and how the electrons are ordered from low energy levels to high energy levels. This is an important topic to keep in mind when teaching physical chemistry.
Not all elements fall under the Aufbau Principle. For example, ruthenium(4d7 5s1), rhodium(4d8 5s1), silver(4d10 5s1), and platinum(4f14 5d9 6s1) are all exceptions to the Aufbau Principle because they are full or semi-filled partial shells.
2. What are the applications of electronic configuration?
The most popular application of electron configuration is the rationalisation of chemical properties in both inorganic and organic chemistry. In fact, electron configuration, along with a simplified form of molecular orbital theory, is a modern equivalent of the concept of valence, which explains the number and types of chemical bonds that an atom can form. To accurately describe a multi-electron system, a large number of electron configurations are required, and energy cannot be allocated to a single configuration. The basic use of electron configuration is the interpretation of atomic spectra.