Structure of the atom : Electronic Configuration, Valence, Atomic Number, Mass number, Isobars, Isotopes


   Atomic Number:

Atomic number (z) is the number of protons in an atom.

Since for a neutral atom, the number of electrons is equal to number of protons, so Atomic Number is also equal to the number of electrons in neutral the atom.

Atomic number = Number of protons = Number of Electrons (for neutral atom)

Example: The atomic number of Iron (Fe) is 26; Fe atom contains 26 protons and 26 electrons.


Mass Number:

The mass number (A) is defined as the sum of the number of protons and neutrons present in the nucleus of an atom. It is also referred as Atomic Mass. 

Example: Mass number of Nitrogen atom is 14 then it contains 7 protons and 7 neutrons.

Mass number = Number of Protons + Number of Neutrons


Notation of element : 

Atomic number Element Mass number  =  Z XA

Example: Sodium atom has notation as 11Na23 .

Question: An atom X has mass number 40 and atomic number 18. Find out the number of protons, number of electrons and number of neutrons present in the atom X?


Number of protons:

Number of protons = Atomic number = 18

Number of electrons:

In an atom number of electrons = Number of protons = 18

Number of Neutrons = Mass number – Atomic number  = 40-18 = 22


Electronic Configuration : 

The arrangement of electrons in the shells is known as electronic configuration.

Electrons are present in fixed energy levels called Shells. These Shells are also called Orbits. These orbits or Shells are represented by the letters K, L, M, N,… or the numbers n = 1, 2, 3, 4,5,.….

Note: Shell and Orbits are not identical ; They are different ideas. Here for simplification, Students are advised to consider the sense same. Shell more correctly indicates a Sphere (a three dimensional object) while Orbits are Circular paths (A two dimensional object). If Bohr’s Model is considered to be valid; Shells and Orbits are similar for indication purpose. 


Rules for accommodating electrons in various shells (Bohr-Bury Rules)

  1. The maximum number of electrons that can be accommodated in any energy level of the atom is given by a formula 2n2 where n is the number of that Orbit number/ Shell Number / Energy level.
Orbit Number  Shell Name Maximum number of Electrons

(Using 2n2)

1 K Shell 2
2 L Shell 8
3 M Shell 18
4 N Shell 32


2. After a series of experiments and a detailed study by scientists like Louis de Broglie, Schrödinger, Somerfield and others proved that shells or energy levels have Sub Shells within them. Electrons are present in these sub shells which constitute a Shell.

Types of Sub Shell : Sub shells are s, p, d and f types.

Every sub-shell can accommodate a fixed number of electrons.

“s” sub shell can hold a maximum of 2 electrons.

“p” sub shell can hold a maximum of 6 electrons.

“d” sub shell can hold a maximum of 10 electrons.

“f” sub shell can hold a maximum of 14 electrons.

Arrangements of Shell and Sub – Shells  : 

Shell Name Sub – Shell name  Number of Electrons

in that sub shell 

K shell (1st Shell) 1s 2
L Shell (2nd Shell) 2s




M Shell (3rd Shell) 3s 






Energy Order of different Sub-Shells : 


3. Electrons are filled up in an atom from Lower Energy Level to Higher Energy Level. This rule is called Aufbau Principle.


Electronic Configuration of Sodium atom: 

Sodium has atomic number 11 and mass number 23.

The nucleus of sodium has 11 protons and 12 neutrons and it is surrounded by 11 electrons.

Now Its electronic configuration is 1s2, 2s2, 2p6, 3s1

Note: you can write electronic configuration easily, follow steps :

1. Write the sub-shell energy sequence:  1s<2s<2p<3s<3p<4s<3d<4p<5s

2. Check how many electrons are present in the atom 

3. Fill these electrons in the sequence till all electrons are not exhausted; remember sub-shell can accommodate maximum number of electrons as s sub-shell = 2 ; p sub-shell =6, d sub-shell = 10 

The electronic configuration of sodium: 1s2, 2s2, 2p6, 3s1 can also be written as 2, 8, 1 (indicating 1st shell has 2 electrons, 2nd shell has 8 electrons, 3rd shell has 1 electron).

K shell or I shell = 2 electrons

L– Shell or II shell = 8 electrons

M– Shell or III shell = 1 electron

∴ Electronic configuration of sodium atom = 2, 8, 1

 Similarly the electronic configuration of other atoms are : 

Element  Atomic Number  Electronic Configuration 

Type 1 

Electronic Configuration 

Type 2

Hydrogen (H) 1 1s1 1
Lithium (Li) 3 1s2, 2s1 2,1
Carbon (C) 6 1s2, 2s2, 2p2 2,4
Aluminium (Al) 13 1s2, 2s2, 2p6,3s2 , 3p1 2,8,3
Scandium (Sc) 21 1s2, 2s2, 2p6, 3s2, 3p6,4s2,3d1 2,8,9,2
Vanadium (V) 23 1s2, 2s2, 2p6, 3s2, 3p6,4s2,3d3 2,8,11,2
Iron (Fe) 26 1s2, 2s2, 2p6, 3s2, 3p6,4s2,3d6 2,8,14,2


Note the changes in the electronic configuration of Sc, V, Fe .

Q. Now try to write the electronic configuration of

(a) Boron (B) atomic number = 5

(b) Florine (F) atomic number = 9

(c) Titanium (Ti) atomic number = 22

(d) Nickel (Ni) atomic number = 28 


Valence electrons: 

The electrons, which are present in the outermost shell of an atom are called valence electrons.

Example: Sodium 

Atomic number of sodium is 11

Electronic configuration of sodium is 2, 8, 1

In sodium 3rd shell is the outermost shell (valence shell). In this shell it has 1 electron. Hence the number of valence electrons present in sodium is 1.

The chemical properties of an atom are dependent on these valence electrons. Since they are ones that are participate in a chemical reaction.

  • Elements having valence electrons 1, 2 or 3 are called metals.

Exception: Hydrogen has 1 valence electron but it is not consider as a metal.

  • These elements lose electrons easily and form a positively charged ion called cation.

Example: Na – e → Na+

  • Elements having valence electrons 4, 5, 6 or 7 are called as Non-metals
  • These elements gain electrons and forms a negatively charged ion called anion.

Example: F + e → F

  • Elements with same number of valence electrons have similar chemical properties; whereas the elements with different valence electrons have different chemical properties.


The combining capacity of the atoms to form molecules either with same or different elements is defined as Valency. 

Valency can be considered as number of Valence electrons; but in some cases it may have different meaning . It is due to the combining property of atoms to form stable molecules. 

The valency of element is either equal to the number of valency electron is it atom or equal to number of electron required to complete to octet (eight electrons) in valency shell.

For example : 

  • Atom contains less than four electrons in its outermost shell, the Valency of an atom is equal to the number of electrons present in the valence shell. More correctly it is called Electrovalency.

Example: Sodium has one electron in its outermost shell, so the Valency of sodium is 1.

Calcium has two electrons in its outermost shell, so the valency of calcium is 2.

Aluminum has three electrons in its outermost shell, so the valency of aluminum is 3.

  • If the outer shell has more than four electrons, the valency = 8 – the number of electrons in the outer shell. More correctly it is called Covalency. 
  • Covalency :Number of electrons shared by one atom of an element to achieve inert gas configuration.


For Chlorine (atomic number = 17) electronic configuration is 1s2, 2s2, 2p6, 3s2, 3p5 or 2,8,7.

Its outermost shell has 7 electrons , therefore its valency can be either 7 or 8-7=1.

1 valency indicates Cl combining capacity is 1 or it can form 1 chemical bond in its molecules.


Octet Rule: 

An element with 8 electrons in an outermost-shell is said to possess a complete Octet.

Atoms or ions with octet configuration are stable.

Example:  Noble gases possess complete Octet configuration ; they are nonreactive and stable (Except Helium, which possess 2 electrons in its outermost shell).

Atomic mass:  The total number of protons and neutrons present in one atom of an element is called atomic mass.

Helium has two protons and two neutrons.

Mass number of He = 2 + 2 = 4

Mass of He atom 2 + 2 = 4u (unified mass)


Isotopes:  Atoms of the same element with the same atomic number but different atomic masses  are called Isotopes.

Examples: 1H1, 1H2, 1H3 are the isotopes of hydrogen,6 C12 ; 6C14  are isotopes of carbon.

The chemical properties of all the isotopes of an element are the same. This is due to the presence of same number of electrons.




Isobars:  Atoms of different elements with different atomic numbers but have the same mass number are called Isobars.

Example: 18 Ar40, 19 K40 and 20Ca40.


Reasons for chemical reactivity of an atom: The chemical activity of an atom depends on the number valence electrons. An atom with complete octet configuration is chemically inert and it does not participate in chemical reactions.

Example: Noble gases.

The atoms of element with incomplete octet configuration are chemically active. These elements combine with others to attain stable electronic configuration (octet configuration). 

In simple words, atoms combine together so that they acquire 8 electrons in their outermost shell or valence shell.

Example 1: Hydrogen has one electron and it requires one more electron to attain the nearest inert gas configuration. To achieve this each hydrogen atom contributes an electron and form hydrogen molecule.


Example 2: Chlorine atom has 7 electrons in its valence shell and it requires 1 more electron to attain the stable octet configuration. To achieve this, each chlorine atom shares an electron with each other and form chlorine molecule.


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