`=>` The mass number (A) is the total number of neutrons and protons present in the nucleus of an atom of an element.
`=>` Mass number (A) = number of protons + number of neutrons = atomic number + number of neutrons
`=>` An atom with atomic number and mass number is represented as `text()_(Z)X^A`.
`=>` Protons and neutrons are collectively called nucleons.
`=>` The number of neutrons in an atom is equal to the difference between the mass number and the atomic number, or (A- Z).

Isotopes : Isotopes were discovered by F. Soddy. Atoms that have the same atomic number but different mass numbers are called isotopes. Isotopes have identical chemical properties but differ in physical properties.

e.g. `text()_1H^1 ` (Protium), `text()_1H^2` (Deuterium) and `text()_1H^3` (Tritium), `text()_8O^(17)` and `text()_8O^(18)`. The element
polonium (Po) possesses maximum isotopes. `text()_1H^1` is the lightest isotope.

Isobars : Isobar is derived from Greek words : 'isos' means equal and 'baros' means weight. Thus, atoms of different elements having same mass number (A) but different atomic number (Z) are termed as isobars. Isobars are the atoms of different elements and they possess different physical and chemical properties.

e.g. `text()_1H^3` and `text()_2He^3 ; text()_(18)Ar^(40) , text()_(19)K^(40)` and `text()_(20)Ca^(40) ; text()_(52)Te^(130) , text()_(56)Ba^(130)` and `text()_(54)Xe^(130)`.

Isotones : The atoms of an element which have atomic numbers and mass number both different but the number of neutrons in atomic nuclei are same called isotones.

e.g. `text()_1H^3` and `text()_2He^4 , text()_(15)P^(31)` and `text()_(16)S^(32) , text()_(19)K^(39)` and `text()_(20)Ca^(40)`.

lsoelectronic : lsoelectronic species have same number of electrons. e.g., `Ne , Na^+ , Mg^(2+)` all have 10 electrons.

Position and nature of electron is completely described by four sets of quantum numbers. The characteristics of each of the quantum numbers are given below :

THE PRINCIPAL QUANTUM NUMBER :

`=>` The principal quantum number n describes the average distance of the orbital from the nucleus — and the energy of the electron in an atom.

`=>` It can have positive integer (whole number) values: 1, 2, 3, 4, and so on. The larger the value of n, the higher the energy and the larger the orbital.

`=>` Chemists sometimes call the orbitals electron shells.

THE ANGULAR MOMENTUM QUANTUM NUMBER :

`=>` The angular momentum quantum number l describes the shape of the orbital, and the shape is limited by the principal quantum number n: The angular momentum quantum number l can have positive integer values from 0 to n–1. For example, if the n value is 3, three values are allowed for l: 0, 1, and 2.

`=>` Orbitals that have the same value of n but different values of l are called subshells. These subshells are given different letters to help chemists distinguish them from each other. The following table shows the letters corresponding to the different values of l.

`=>` When chemists describe one particular subshell in an atom, they can use both the n value and the subshell letter — 2p, 3d, and so on. Normally, a subshell value of 4 is the largest needed to describe a particular subshell. If chemists ever need a larger value, they can create subshell numbers and letters.

THE MAGNETIC QUANTUM NUMBER :

`=>` The magnetic quantum number is designated as:

`m_1`

`=>` This number describes how the various orbitals are oriented in space. The value of this number depends on the value of l. The values allowed are integers from –l to 0 to +l. For example, if the value of l = 1 (p orbital), you can write three values for this number: –1, 0, and +1. This means that there are three different p subshells for a particular orbital. The subshells have the same energy but different orientations in space.

`=>` The second row (b) of the figure shows how the p orbitals are oriented in space. Notice that the three p orbitals correspond to magnetic quantum number values of –1, 0, and +1, oriented along the x, y, and z axes.

THE SPIN QUANTUM NUMBER :

`=>` The fourth and final quantum number is the spin quantum number, designated as: `m_s`

`=>` This number describes the direction the electron is spinning in a magnetic field — either clockwise or counterclockwise.

`=>` Only two values are allowed: +1/2 or –1/2. For each subshell, there can be only two electrons, one with a spin of +1/2 and another with a spin of –1/2.

The arrangement of electrons in various shells, subshells or orbitals of an atom is known as electronic configuration of the element. Filling of electrons in different orbitals is governed by the following rules.

`=>` According to `(n +l)` rule, the lower the value of `(n +l)` for an orbital the lower is its energy e.g. between `3d` and `4s`, the `4s (4 + 0 = 4)` be filled before `3d (3 + 2 = 5)`.

`=>` If two orbitals have same value of `(n + l)`, the orbital with lower value of `n` will be filled first e.g. between `2p` and `3s`, `2p (2+ 1 = 3)` will be filled first than `3s ( 3 + 0 = 3)`. The order of increasing energies is summed as `1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d,6p,7s,5f,6d,7p`

`text(Note) :` `text()_(24)Cr` and `text()_(29)Cu` do not obey this law.

`text()_(24)Cr = 1s^2 , 2s^2 , 2p^6 , 3s^2 , 3p^6 , 3d^5 , 4s^1`

`text()_(29)Cu = 1s^2 , 2s^2 , 2p^6 , 3s^2 , 3p^6 , 3d^(10) , 4s^1`.

`=>` The completely filled and half-filled sub-shells have lesser energy and thus assume more stability than any other arrangement. Thus `3d^5 , 4s^1` and `3d^(10) , 4s^1` are more stable arrangement than `3d^(4) , 4s^2` and `3d^9 , 4s^2` respectively.

According to this rule, "Pairing of electrons in a sub-shell starts after all the available atomic orbitals or the sub-shell are singly filled (half-filled)."