NCERT Solutions Class 9th Science Chapter – 3 Atoms And Molecules Notes

April 1, 2022

NCERT Solution Class 9th Science Chapter – 3 Atoms And Molecules

Textbook	NCERT
Class	9^th
Subject	Science
Chapter	3^rd
Chapter Name	Atoms And Molecules
Category	Class 9^th Science
Medium	English
Source	Last Doubt

NCERT Solutions Class 9th Science Chapter – 3 Atoms And Molecules Notes What is the difference atom and molecule,What is atom and molecule with example,How many atoms in a molecule, What are 3 differences between atoms and molecules, What are 5 examples of molecules,What is called atom,What are 3 examples of atoms,What is a molecule example,Is water an atom,What color is an atom,What is atom made of molecule.

NCERT Solution Class 9th Science Chapter – 3 Atoms And Molecules

Chapter – 3

Atoms And Molecules

Notes

Laws of Chemical Combination

• Law of Conservation of Mass

• Law of Constant Proportion

Law of Conservation of Mass

• Mass Can neither be created nor be destroyed.

• eg. C+O₂→CO_{2 12gm→ 44gm}

Law of Constant Proportion

• Compound Consists of the elements combined together in the same ratio.

• H2O=2:16 1:8

Daltons Atomic Theory

• (Atom) are tiny particle
• Indivisible (can neither be created nor destroyed)
• Identical mass and Chemical Properties (same element)
• Different mass and Chemical Properties (diff. elements)
• Combine in the same ratio

Atom

• Smallest particle of an element
• Each atom shows all the properties of element
• Form Ions

Atomic mass

• Mass of an atom
• equals to 1/12 of mass C-12 atom

Existence of Atom

• Reactive Molecule, Ions
• Free
• Molecule can be
• Monoatomic (free)
• Polyatomic Molecule (Reactive)

Ions

• Cations (+vly Charged)
• Anions (-vly Charged)

Molecule

• Smallest particle of elements or compounds
• Properties of the substances are the properties of molecules

Types of Molecule

• Homoatomic
• Heteroatomic

Molecular mass

Sum of masses of all atoms.

Formula unit mass

Sum of atomic mass of ions and atoms.

Atomicity

No. of atoms present in one molecule.

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.

Law of Chemical Combination

• Law of Conservation of Mass
• Law of constant Proportion
• Law of Multiple Proportion
• Law of Reciprocal Proportion
• Gay Lussac’s law of combining Volume

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.”

For example, A + B → AB

Then, Reactant Product

where, m_A + m_B = m_AB

m_A = Mass of A

m_{B = Mass of B}

m_AB = Mass of AB

2H₂ (g) + O₂ (g) → 2H2O (1)

2×2=4gm 2×16=32gm 2×(2+16)=36gm

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

Solution

Sodium Carbonate + Ethanoic acid → Sodium ethanoate + CO₂ + H₂O

↓	↓
Reactants	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:
5.3gm + 6.0gm = 8.2gm + 2.2gm + 0.9gm
Or 11.3gm = 11.3gm
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., m/mo-2/16 = 1/8

• 36 gm of H₂O⇒ 32 gm of oxygen + 4 gm of hydrogen,
i.e., m/m 4/32 = 1/8

• 09 gm pf H₂O⇒ 08 gm of oxygen + 1 gm of hydrogen,
i.e., m/m = 1/8

From the above three cases, differently weighing H, O 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 ‘H, O’ 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?

Solutin

m_H/m_O=1/8 Given in equation (For H₂O)

But, m_H=3.0gm (given)
Or 3/m_O = 1/8
Or m_O=24gm

.: 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 as follows:

• 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.

Atom

• 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 of Hydrogen is 0.37 x 10-10 m
or 0.037 nm.

Such that, 1 nm = 10^-9 m

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

Elements	Symbol
Aluminium	Al
Argon	Ar
Barium	Ba
Calcium	Ca
Carbon	C
Chlorine	Cl
Cobalt	Co
Copper	Cu
Fluorine	F
Gold	Au
Hydrogen	H
Iodine	I
Iron	Fe
Lead	Pb
Nitrogen	N
Oxygen	O
Potassium	K
Silicon	Si
Silver	Ag

Sulphur	S
Zinc	Zn

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 respectively.

Atomic Mass Unit

This means atomic mass unit 1/12th of Carbon – 12

Atomic Mass of some elements

Element	Symbol	Atomic Mass	Element	Symbol	Atomic Mass
Hydrogen	H	1u	Sodium	Na	23u
Helium	He	4u	Magnesium	Mg	24u
Lithium	Li	7u	Aluminium	Al	27u
Beryllium	Be	9u	Silicon	Si	28u
Boron	B	11u	Phosphorous	p	31u
Carbon	C	12u	Sulphur	s	32u
Nitrogen	N	14u	Chlorine	Cl	35u
Oxygen	O	16u	Potassium	k	39u
Fluorine	F	19u	Calcium	Ca	40u
Neon	Ne	20u	Iron	Fe	56

How do atoms exist?

Molecule

• 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.

E.g., ‘H, O’ 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 (containing atom of same element) and
Heteroatomic molecules or compounds (containing atoms of different elements)

Molecules
Homoatomic Molecules of Elements	Heteroatomic molecules Compounds
Eg, O₂,N₂,O₃,S₈,P₄ etc.	Eg, H₂O, NO₂, SO₂ etc.

Molecules of Some Compounds
Compound	Combining Elements
Water (H₂O)	Hydrogen, Oxygen
Ammonia (NH₃)	Nitrogen, Hydrogen
Carbon Dioxide (CO₂)	Carbon, Oxygen
Hydrogen Chloride (HCl)	Hydrogen, Chloride
Methane (CH₄)	Carbon, Hydrogen
Ethane (C₂H₆)	Carbon, Hydrogen
Sodium chloride (NaCl)	Sodium, chlorine
Copper Oxide (CuO)	Copper and Oxygen

Atomicity

The number of atoms present in one molecule of an element is called its atomicity.

Name	Formula	Atomicity
1. Argon	Ar	Monoatomic (1) Noble gasses constitute
2. Helium	He	Monoatomic (1) Monoatomic molecules
3. Oxygen	O₂	Diatomic (2)
4. Hydrogen	H₂	Diatomic (2)
5. Phosphorus	P₄	Tetratomic (4)
6. Sulphur	S₈	Polyatomic (8)
7. Ozone	O₃	Triatomic (3)

Chemical formulae

It is the Symbolic representation of the composition of 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. E.g., CaO,NaCl,CuO.
• When polyatomic ions are used,the ions are enclosed in brackets before writing the number to show the ratio. e.g., Ca(OH)₂,(NH₄)2SO₄

Molecular Mass

It is the sum of atomic masses of all the atoms in a molecule of that substance.

e.g., Molecular mass of H₂O = 2 x Atomic mass of Hydrogen + 1 x
Atomic mass of Oxygen

So, Molecular mass of H₂O= 2 x 1+1 x 16 = 18 u

Formula Unit Mass

It is the sum of atomic mass of ions and atoms present in formula for a compound.

e.g., In NaCl, Na = 23 a.m.u. , C1 = 35.5 a.m.u.
So, Formula unit mass = 1 x 23 + 1 x 35.5 = 58.5 u

Ions

An ion may be defined as an atom or group of atoms having positive or negative charge.

Some Positively Charged Ions : Na⁺,K⁺,Ca2⁺,Al3⁺ etc.

Some Positively Charged Ions

Cl^– (Chloride ion), S^2- (Sulphide ion), OH^– (Hydroxide ion), SO₄^2- (Sulphate ion) No₃ (Nitrate ion)

Ions

Monoatomic Ions	Polyatomic Ions
Mg²⁺ (Magnesium ion)	NH₄⁺ (Ammonium ion)
Na⁺ (Sodium ion)	CO₃^2- (Carbonate ion)
CI^– (Chloride ion)	SO₄^2- (Sulphate ion)
Al³⁺ (Aluminium ion)	OH^– (Hydroxide ion)

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.

Example

(a) Atomic mass of hydrogen (H) is 1 u. Its molar mass is 1 g/mol.

(b) Atomic mass of nitrogen is 14 u. So, molar mass of nitrogen (N) is 14 g/mol.

(c) Molar mass of S = Mass of S x 8 = 32 x 8 = 256 g/mol

(d) 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 1033 Particles (atoms, molecules or ions) of a substance is called a mole of that substance.

1 mole of atoms = 6.022 x 10²³ atoms

1 mole of molecules = 6.022 x 10²³ molecules

Example, 1 mole of oxygen atoms = 6.022 x 10²³ oxygen atoms

6.022 x 10²³ is Avogadro Number

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

(i) Number of moles (n) = Given number of particlas \Avogadr0’s number

(ii) Number of moles (n) = n=N\N₀m= MxN\N₀

(iii) m\M= N\N₀

(iv) Percentage of any atom in given compound = Mass of element x100 / Mass of compound

Example

Calculate no. of iron atoms in a piece of iron weighing 2.8 gm (Atomic mass 54 u).

Solution : 1 mole of iron = 56 gm (Gram atomic mass of iron)
1 mole of iron element contains 6.022 x 10²³ atoms of iron.
So, 56 gm of iron = 6.022 x 10²³ atoms
2.8 gm of iron = 6.022/56 x 102 x 2.8
= 3.011 x 1022 atoms

Example

Mass of one molecule of a substance is 5.32 x 1023 g. What is its molecular mass?

Solution : Mass of 1 molecule of substance
= 5.32 x 1023 g
Mass of 6.022 x 1023 molecules of substance
= 5.32 x 1023 x 6.022 x 1023
= 32g

Example

Calculate the mass of 0.5 mole of N, gas.

Solution : 1 mole of N₂ = Gram molecular mass of N₂
Or 1 mole of N₂ = 28 gm
0.5 mole of N₂ gas = 0.5 x 28
= 14 gm of N₂

Example

Calculate the total number of O, molecules present in 8 gm of O₂

Solution: Gram molecular mass of O₂
= 6.022 x 10²³ 0₂ molecules
Or 32 gm of O₂ = 6.022 x 10²³ O₂ molecules
Or 8 gm of O₂ = 6.022 x 10²³ x 8/32 O₂ molecules
= 1.51 x 103 O₂ molecules

Question 1. Give one example of cation and anion.

Cation => Na+
Anion => Cl–

Question 2. Give one difference between cation and anion.

Cations are positively charged ion.
Anions are negatively charged ion.

Question 3. Give the chemical formula for ammonium sulphate.

Ammonium sulphate
NH4+ SO42-
Chemical formula —-> (NH4)2S04.

Question 4. Define atom.

The smallest particle of matter, which can take part in a chemical reaction is called atom.

Question 5. How do atoms exist?

Atoms exist in the form of atom, molecule or ions.

Question 6. Define atomicity.

The number of atoms constituting a molecule is known as its atomicity.

Question 7. Define molecule.

The smallest particle of an element or compound which can exist independently is called molecule.

Question 8. Who coined the term atom?

John Dalton coined the term atom.

Question 9. Give the atomicity of phosphorous and nitrogen.

The atomicity of phosphorus is P4 i.e., 4.
The atomicity of nitrogen is N2 i.e., 2.

Question 10. What is an ion?

Charged atom is called as an ion. The ion can be positively charged called cation or negatively charged called anion.

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