Atomic Structure

A tiny unit of every known matter that ends up as a chemical element comes under the term of an atom. It has a size of around 100 picometres. Most of the atom’s mass is because of the nucleus. An atom is made up of a nucleus (in the middle or centre) that has neutrons and protons, and the electrons, which are negative in charge, revolve around the nucleus. 

Due to the nature of quantum mechanics, no image is completely capable of predicting the different properties of an atom, and physicists must have used complementary or different images of the atom to explain their different properties. The electrons in an atom behave like particles when orbiting the nucleus. Elsewhere, the electron also behaves like a moving wave around the nucleus. Such a wave pattern, referred to as an orbit, represents the distribution of individual electrons. The behaviour or nature of atoms is strongly affected by these orbital properties, the chemistry of which can be determined by a group of orbitals called shells.

In the absence of charge, neutrons are not repelled by electron clouds or nuclei, making them a useful tool for studying the structure of atoms. Because of the presence of different counts of subunit particles in atoms, they have distinct properties for different atoms. 

This is the reason for the unique properties of the various elements. 

Basic terminologies of atomic structure

1. Atomic number: This is one of the most important properties of an atom (usually indicated by the letter Z). It can be expressed as the number of units of positive charge or proton in the nucleus. 
2. Atomic mass: Neutrons are the key reason for the mass of the atom and can affect the same by changing its count, but not its chemistry. Therefore, a nucleus with 4 protons and 4 neutrons has the same chemistry as a nucleus with 4 protons and 8 neutrons, although the two masses are different. 
3. Electron: J.J. Thomson first discovered the presence of electrons in the nucleus in 1897, marking the beginning of modern science. Many of the properties of atoms depend upon the number and arrangements of electrons present. The mass of an electron is 9.1*10-31kg. 
4. Proton: Proton was first discovered by Sir Ernest Rutherford in 1919. It has the same charge but is opposite to an electron. The mass of a proton is 1.67*10-27kg.
5. Neutron: Neutrons are another type of particle found in the nucleus and were discovered by Sir James Chadwick, a British physicist. Neutrons have no charge and have the same mass as protons. 
6. Nucleus: The proton and neutron, which have similar mass and are substantially more heavy than the electron, are the nucleus’ fundamental constituents. Both have magnetic fields that are inherent to them. A more accurate representation of the nucleus would be a boiling cauldron filled with hundreds of various sorts of particles swarming around protons and neutrons.
7. Orbits and energy levels: Electrons can’t exist at any random distance from the nucleus; they can only dwell in certain regions known as permitted orbits. The angular momentum of an electron in orbit is necessary, like everything else in the quantum world.
8. Electron shells: A shell is made up of all orbitals with the same n value. Subshells corresponding to varying rates of spin and orientation of orbitals, as well as electron spin orientations, may exist inside each shell. In general, the further a shell is from the nucleus, the more subshells it has.
9. Atomic orbitals: Electrons are filled in the shell and subshell levels in a semi-regular manner. Electrons travel into the second-level s subshell and then into the p subshell after filling the first shell level (with just a subshell). 
10. Atomic bonds: The outer electrons of atoms can form bonds in three different ways. It is possible to transfer electrons from one atom to another. Electrons can be shared between atoms in close proximity. 

Electronic configuration of an atom

Different principles are used to fill up the electron in the s,p,d and f orbitals. 

1. Aufbau’s principle: Electrons should be filled in the increasing manner of energy of orbitals: The lower energy levels should be filled first, followed by the higher energy levels. The desired order of energy is 1s, 2s, 2p, 3s, 3p, 4s, 3d.
2. Pauli’s exclusion: It states that two electrons cannot have the same four quantum numbers, electrons having opposite spin can be placed in the same energy state.
3. Hund’s rule: This rule states that when filling identical energy orbitals called degenerate orbital, all degenerate orbitals must be first singly filled, followed by pairing.

Conclusion 

Every unit of any ordinary matter that ends up as a chemical element is termed an atom. It has a size of around 100 picometres. Most of the atom’s mass is because of the nucleus. Atoms are made up of a nucleus that has protons and neutrons, the electrons are negative in charge and revolve around the nucleus. Because of the presence of different counts of subunit particles in atoms, they have distinct properties for different atoms. This is the reason for the unique properties of various elements. 

Three rules are used for filling electrons in orbitals: Aufbau’s principle, Pauli’s exclusion principle and Hund’s rule.