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# Distribution of Electrons in Different Orbits/Shells

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Last updated date: 07th Sep 2024
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## What is Electronic Configuration?

The distribution of electrons in shells of the atomic orbitals of an atom is known as its electronic configuration. An atom consists of subatomic particles like protons, neutrons, and electrons. Protons and neutrons are present in the nucleus of an atom, while electrons are the ones that revolve around the nucleus of an atom in fixed orbits, which is explained by Bohr’s atomic model.

An atomic number of an element is equal to the number of electrons or protons (in a neutral atom, the number of electrons is equal to the number of protons) present in an atom. Now there’s a particular way these electrons are arranged in different orbitals, called the electronic configuration of an atom.

## Bohr’s Atomic Model - Salient Features

The Bohr Atomic model states that a small, positively charged nucleus is surrounded by revolving, negatively charged electrons in set orbits. He concluded that an electron would have more energy if it is farther from the nucleus, whereas an electron would have less energy if it is closer to the nucleus. The postulates are as follows -

1. Electrons are negatively charged and revolve in a fixed circular path around the nucleus known as shells/orbits/energy level.

2. Each of these orbits/shells have a fixed energy level.

3. Integers such as n=1,2,3, and so on are used to represent the various energy levels. These are referred to as principle quantum numbers. The range of quantum numbers may change from the lowest energy level (nucleus side n=1) to the greatest energy level.

4. There are two methods to express the various energy levels or orbits, such as 1, 2, 3, 4..., and their corresponding shells/orbits are K, L, M, and N... shells. The term "ground state" refers to the electron's lowest energy level.

5. When an electron in an atom needs to go from a lower to a higher energy level, it gains the necessary energy, and when it needs to move from a higher to a lower energy level, it loses energy.

## How are Electrons Distributed in Different Orbits/Shells?

As we know, outside of an atom's nucleus, a cloud of negatively charged subatomic particles called electrons are distributed in different shells. The potential energy in various orbits determines their arrangement. The different energy levels are referred to as 1, 2, 3, 4… and the equivalent shells are referred to as K, L, M, N, and so on.

Starting in the centre, the shells progressively spread outward. K-shell will have the least energy. The L shell will have more energy than the K shell since it is somewhat further from the nucleus. The energy level will be highest in the outermost shell. Any atom will be most stable when it has the least energy.

To reach the state of least energy, an atom must first occupy the lowest energy level. The higher energy levels will eventually fill with electrons. As a result, electrons will begin to fill the K shell before moving on to the L, M, N, and so on.

## Bohr-Bury Scheme

The Bohr-Bury Scheme determines how the electrons are distributed throughout the orbits or shells of the atoms. The concept behind the Bohr-Bury Scheme is as follows:

• The electrons in an atom are grouped in several shells around the nucleus, and they first inhabit the shell closest to the nucleus since it has the lowest energy.

• The first and innermost shell is known as the K-shell, which may accommodate up to two electrons.

• The L-shell, the second shell, may hold up to eight electrons.

• A total of eighteen electrons can fit in the third shell, known as the M-shell.

• The total or the maximum number of electrons in each shell can be given by the formula 2n2, where n (n=1,2,3,4….) is the shell number or the principal quantum number.

## Filling Up of Atomic Orbitals

It is governed by 3 principles, namely:

• Aufbau Principle: According to it, electrons occupy atomic orbitals in ascending order of orbital energy level, meaning that lower energy levels are occupied before higher energy levels.

• Pauli’s Exclusion Principle: Pauli's Exclusion Principle states that no two electrons in the same atom may have the same values for all four of their quantum numbers.

• Hund’s Rule: Before becoming doubly occupied, every orbital in a subshell is solely filled by one electron, and all of the electrons in singly occupied orbitals have the same spin.

### 1. Helium

Orbital diagram of Helium

Atomic number = 2

Number of electrons = 2

Electronic configuration = 2

The K shell can occupy a maximum of 2 electrons hence only 1 Shell, which is the K shell, is required to distribute the 2 electrons of Helium.

### 2. Oxygen

Orbital Diagram of Oxygen

Atomic number = 8

Number of electrons = 8

Electronic configuration = 2,6

K shell  (1st orbit) may hold a maximum of 2 electrons. The remaining electrons (6 electrons) will fit into the second orbit, which is the L shell and can accommodate a maximum of 8 electrons. Hence, Oxygen requires a total of 2 shells to distribute its electrons.

## Important Questions

1. What is electronic configuration?

Ans: The arrangement or distribution of electrons in the shells of an atom is defined as the electronic configuration of an element.

2. What are shells/orbits in an atom?

Ans: Electrons revolve around the nucleus in a fixed circular path known as shells/orbits. There are K, L, M, N…..shells, and these can also be represented as integers based on the principal quantum number.

## Summary

Subatomic particles called electrons are negatively charged and distributed outside the atom's nucleus in fixed orbits/shells. Based on their potential energies in various orbits, the arrangement is determined. Electronic configuration describes how the electrons are distributed throughout the energy shells.

It is based on the Bohr-Bury scheme, which states that the maximum number of electrons that may be found in a specific energy shell of an atom is determined by the formula 2n2, where n is the number of energy shells. 1, 2, 3, and 4 are the numbers that represent the various energy levels. The designations for the corresponding shells are K, L, M, N, and so on.

## Practice Questions

1. How many shells are required to distribute the 17 electrons found in chlorine atoms?

1. 1

2. 2

3. 3

4. 4

2. Which of the following principles govern the filling up of the atomic orbitals?

1. Aufbau principle

2. Pauli’s exclusion principle

3. Hund’s rule

4. All of the above

1. (c)

2. (d)

Competitive Exams after 12th Science

## FAQs on Distribution of Electrons in Different Orbits/Shells

1. What is a principal quantum number?

The letter "n" represents principal quantum numbers. They identify the atom's outermost shell of electrons. Since it describes the most probable distance between the nucleus and the electrons, a higher value of the principal quantum number denotes a farther separation between the electron and the nucleus (which, in turn, implies a greater atomic size). The value of the principal quantum number can be any integer with a positive value equal to or greater than one.

2. Mention a few uses of electronic configuration.

Understanding electronic configuration is a crucial and fundamental aspect of Chemistry. It serves as the periodic table's foundation. Any element's electronic configuration will also have an impact on the stability of an atomic orbital. Additionally, it aids in our comprehension of how elements are organised into various periods and groups in the periodic table. It also aids in the comprehension and understanding of the chemical bonding present in each element. It discusses the many characteristics and unusual characteristics of individual elements.

3. What are the limitations of Bohr’s Atomic Model?

Bohr’s atomic model violates Heisenberg's Uncertainty Principle. Heisenberg asserts that an object can't have a known location and momentum simultaneously, yet the Bohr atomic model theory assumes that electrons do.

• The Bohr atomic model theory yields inaccurate spectrum predictions when bigger atoms are considered. This is true for smaller atoms like hydrogen.

• It was unable to explain the Zeeman effect, which occurs when a magnetic field causes the spectral line to split into numerous components.

• It could not explain the Stark effect, which occurs when an electric field causes the spectral line to split into smaller lines.