Class 9 Science Chapter 3 Atoms and Molecules NCERT Notes

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Chapter 3 Atoms and Molecules Class 9 Science CBSE NCERT Notes

Ancient Indian and Greek philosopher have wondered about unknown and unseen form of matter. Around 500 BC, An Indian Philosopher Maharishi Kanad postulated that if we go on dividing matter (padarth) we shall get smaller and smaller particles. Ultimately, a stage will come across the smallest particles beyond which further division will not be possible. He named these particles Parmanu.

Another philosopher Pakudha Katyayama, elaborated this doctrine and said that these particles exist in combined form. Around the same era, ancient Greek philosophers, Democritus and Leucippus suggestion the same and coined the term atom.

Laws of Chemical Combination

The chemical reaction between two or more substances giving rise to products is governed by certain laws. These laws are called Laws of Chemical Combination.

There are two laws of chemical combinations given after much experimentations by Antoine L. Lavoiser and Joseph L. Proust. Two laws are:

  • Law of conservation of masses
  • Law of constant proportion

Law of Conservation of Mass

According to this law, “Mass can neither be created nor destroyed.”

In a chemical reaction, this law can be understood in the following way: During a chemical reaction total mass of reactants will be equal to total mass of products.”

Example: In a reaction 5.3 gm of sodium carbonate reacted with 6 gm of ethanoic acid. The products were 2.2 gm of CO2, 0.9 gm of H2O and 8.2 gm of sodium ethanoate. Show that these observation are all in agreement with law of conservation of mass.
Sodium carbonate + Ethanoic acid → Sodium ethanoate + CO2 + H2O


Sodium carbonate + Ethanoic acid (Reactants) → Sodium ethanoate + CO2 + H2O (Products)
Now, according to the law of conservation of mass :
Mass of sodium carbonate + Mass of ethanoic acid = Mass of sodium ethanoate + Mass of CO2 + Mass of H2O

Putting values of masses from the equation :
5.3 gm +6.0 gm = 8.2 gm + 2.2 gm +0.9 gm
Or, 11.3 gm = 11.3 gm
Since, LHS = RHS
∴ Law of conservation of mass is in agreement with the given values in equation.

Law of Constant Proportion

According to this law, “A pure chemical compound always contain the same elements combined together in the same proportion by mass irrespective of the fact from where the sample has been taken or from which procedure has it been produced.”

For example:

18 gm of H2O ⇒ 16 gm of oxygen + 2 gm of hydrogen,
i.e., mH/mO = 2/16 = 1/8
36 gm of H2O ⇒ 32 gm of oxygen + 4 gm of hydrogen,
i.e., mH/mO = 4/32 = 1/8
09 gm of H2O ⇒ 08 gm of oxygen + 1 gm of hydrogen,
i.e., mH/mO = 1/8

From the above three cases, differently weighing H2O samples were taken but the ratio of masses of ‘H’ to mass of ‘O’ comes out to be ‘1/8’ is same, proving law of constant proportion.

Likewise, if a sample of ‘H2O’ was taken from anywhere i.e., from well, pond, lake or anywhere the ratio of masses of ‘H’ to ‘O’ will come out to be same as ‘1/8’.

Example: Hydrogen and oxygen combine in the ratio 1:8 by mass to form water. What mass of oxygen gas would be required to react completely with 3.0 gm of hydrogen gas?


mH/mO = 1/8 Given in equation (For H2O)
But, mH = 3.0 gm (given)
or, 3/mO = 1/8
or, mO = 24 gm
∴ Mass of oxygen will be 24 gm.
Or, it will be a sample of 27 gm of H2O where 3 gm of hydrogen is present with 24 gm of oxygen.

Dalton’s Atomic Theory

Based upon laws of chemical combination, Dalton’s Atomic Theory provided an explanation for the Law of Conservation of Mass and Law of Constant Composition.

Postulates of Dalton’s atomic theory are:

  • All matter is made up of very tiny particles called ‘Atoms’.
  • Atom are indivisible particles, which can’t be created or destroyed in a chemical reaction. (Proves ‘Law of Conservation of Mass’)
  • Atoms of an element have identical mass and chemical properties.
  • Atoms of different elements have different mass and chemical properties. Atom combine in the ratio of small whole numbers to form compounds. (proves ‘Law of Constant Proportion’)
  • The relative number and kinds of atoms are constant in a given compound.


According to modern atomic theory, an atom is the smallest particle of an element which takes part in chemical reaction such that during the chemical reaction, the atom maintain its identity, throughout the chemical or physical change.

Atoms are very small and hence can’t be seen even through very powerful microscope.

Atomic radius of smallest atom in hydrogen is 0.37 x 10-10 m or 0.037 nm.

Such that, 1nm = 10-9 m.

IUPAC (International Union of Pure & Applied Chemistry) Symbols of Atoms of Different Elements

Jöns Jacob Berzelius suggested that the symbols of elements be made from one or two letters of the name of the element.

In beginning, the names of element were derived from the name of the place where they were found for the first time. For example name copper was taken from Cyprus.

Some names were taken from specific colours. For example name gold was taken from English word meaning yellow.

Now, IUPAC (International Union of Pure and Applied Chemistry) is an international organization which approves names of element, symbol and units.

Many are first one or two letter of the elements name in English. The first letter of the symbol is allways written in Uppercase and second letter as lowercase. Like, For Hydrogen- H, Aluminum – Al (not AL) etc.

Symbols of some elements are formed from the first letter of the name and the letter, appearing later in the name. Like Chlorine- Cl, Zinc – Zn.

Other symbols have been taken from the name in Latin, German or Greek. For example, the symbol of iron is Fe from its Latin word ferrum. Sodium is Na from Natrium, Potassium is K from Kalium. Gold is Au from Aurum etc.

Atomic Mass

The mass of an atom of an element is called its atomic mass.

In 1961, IUPAC have accepted ‘atomic mass unit’ (u) to express atomic and molecular mass of elements and compounds.

Atomic Mass Unit

The atomic mass unit is defined as the quantity of mass equal to 1/12 of mass of an atom of carbon-12.

1 amu or u = 1/12 ✕ Mass of an atom of C
1 u = 1.66 ✕ 10-27 kg
This means atomic mass unit 1/12th of Carbon-12.

How do Atoms Exist?

Atoms of most elements are not able to exist independently. Atoms form molecules and ions. These molecules or ions aggregate in large number to form the matter that we can see, feel or touch.

Only the atoms of noble gases (such as He, Ne, Ar, Kr, Xe and Rn) are chemically unreactive and can exist in the free state as single atom.


A molecule is a group of two or more atoms which are chemically bonded with each other.

A molecule is the smallest particle of matter (except element) which is capable of an independent existence and show all properties of that substance. For example, ‘H2O’ is the smallest particle of water which shows all the properties of water.

A molecule may have atom of same or different elements, depending upon this, molecule can be categorized into two categories :

Homoatomic molecules are containing atom of same element. For example, O2, N2, O3 etc.
Heteroatomic molecules or compounds are containing atoms of different elements. For example, H2O, NO2, SO2 etc.


Compound composed of metals and non-metals contain charged species. The charged species are known as ions.

Ions may consist of a single atom or a group of atoms that have a net charge on them. An ions can be negatively or positively charged.

A negatively charged ion is called ‘anion’. The positively charged ion is called ‘cation’. For example, Sodium chloride (NaCl) is constituted of positively charged particles of sodium ion (Na+) and negatively charged chloride ions (Cl).

Group of atoms carrying a charge is known as polyatomic ion. Ions can be separated using electrolysis.

Chemical formulae

It is the symbolic representation of the composition of a compound. Characteristics of chemical formulae. The valencies or charges on ion must balance.

When a compound is formed of metal and non-metal, symbol of metal comes first. For example, CaO, NaCl, CuO.

When polyatomic ions are used, the ions are enclosed in brackets before writing the number to show the ratio. For example, Ca(OH)2, (NH4)2SO4

Molecular Mass

It is the sum of atomic masses of all the atoms in a molecule of that substance. For example, Molecular mass of H2O= 2 ✕ Atomic mass of Hydrogen + 1 ✕ Atomic mass of Oxygen.
So, Molecular mass of H2O = 2 ✕ 1 + 1 ✕ 16 = 18 u

Formula Unit Mass

It is the sum of atomic mass of ions and atoms present in formula for a compound. For example, In NaCl, Na = 23 a.m.u. Cl = 35.5 a.m.u.
So, Formula unit mass = 1 ✕ 23 + 1 ✕ 35.5 = 58.5 u

Formulae of simple compounds

Formula of hydrogen chloride

Formula of hydrogen sulphide

Formula of carbon tetrachloride

Formula of magnesium chloride

Formula for aluminium oxide

Molar Mass

The molar mass of a substance is the mass of 1 mole of that substance. It is equal to the 6.022 x 1023 atoms of that element/substance.


Atomic mass of hydrogen (H) is 1 u. Its molar mass is 1 g/mol.
Atomic mass of nitrogen is 14 u. So, molar mass of nitrogen (N) is 14 g/mol.
Molar mass of S = Mass of S ✕ 8 = 32 ✕ 8 = 256 g/mol
Molar mass of HCl = Mass of H + Mass of Cl
= 1= 35.5 = 36.5 g/mol

Mole concept

A group of 6.022 x 1023 Particles (atoms, molecules or ions) of a substance is called a mole of that substance.

1 mole of atoms = 6.022 ✕ 1023 atoms
1 mole of molecules = 6.022 ✕ 1023 molecules
Example, 1 mole of oxygen = 6.022 ✕ 1023 oxygen atoms

6.022 ✕ 1023 is Avogadro number.

1 mole of atoms of an element has a mass equal to gram atomic mass of the element.

Important Formulae

(i) Number of moles (n) = Given mass/Molar mass = m/M
(ii) Number of moles (n) = Given number of particles/Avogadro’s number
n = N/N0
(iii) m/M = N/N0
m = M✕N/N0
(iv) Percentage of any atom in given compound = Mass of element ✕ 100/Mass of compound

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