Class 9 Chapter 3 Long Qa Science Cbse Ncert

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Extra 20 important long questions and answers containing from the chapter:3 "Atoms and Molecules" for Class 9 CBSE Science:-

1. Explain the Law of Conservation of Mass with an experiment.

Answer:
The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction.

Experimental Verification:

Take

barium chloride (

BaCl

₂) solution

and

sodium

sulphate

(

Na₂SO

₄) solution

in two separate test tubes.

Weigh both solutions separately.

Mix the two solutions in a conical flask.

A white precipitate of

barium

sulphate

(

BaSO

₄)

forms.

Weigh the flask after the reaction.

Observation:

The total mass before and after the reaction remains the same, proving the law.

2. State and explain the Law of Constant Proportions with an example.

Answer:
The Law of Constant Proportions states that a given compound always contains the same elements in a fixed proportion by mass, irrespective of its source or method of preparation.

Example:

Water (H₂O) always contains hydrogen and oxygen in a

fixed ratio of 1:8 by mass

If we take 9g of water, it contains

1g of hydrogen

and

8g of oxygen

, proving the law.

3. Describe Dalton’s Atomic Theory with its postulates.

Answer:
John Dalton proposed the Atomic Theory in 1808. Its main postulates are:

All matter is made of tiny indivisible particles called atoms.

Atoms of a given element are identical

in size, mass, and properties.

Atoms of different elements are different

in mass and properties.

Atoms combine in whole numbers to form compounds.

Chemical reactions involve rearrangement of atoms; atoms are neither created nor destroyed.

Limitations:

It failed to explain

isotopes and subatomic particles

(electrons, protons, neutrons).

4. Define atomic mass and molecular mass. How are they different?

Answer:

Atomic Mass:

The mass of an atom compared to

1/12th the mass of a carbon-12 atom

Example: Hydrogen =

amu

, Oxygen =

amu

Molecular Mass:

The sum of atomic masses of all atoms in a molecule.

Example: CO₂ =

12 + (16 × 2) = 44

amu

Difference:

Feature	Atomic Mass	Molecular Mass
Definition	Mass of a single atom	Sum of atomic masses in a molecule
Example	Oxygen = 16 amu	CO₂ = 44 amu

5. What is a mole? Explain its significance with examples.

Answer:
A mole is a unit that represents 6.022 × 10²³ particles (atoms, molecules, or ions) of a substance.

Significance:

Converts atomic/molecular mass into grams

Used to count microscopic particles

(atoms/molecules).

Relates mass, volume, and number of particles

Example:

1 mole of

H₂O = 6.022 × 10²³ molecules

1 mole of

O₂ = 32 g

6. Differentiate between an atom and a molecule.

Answer:

Feature	Atom	Molecule
Definition	Smallest unit of an element	Smallest unit of a compound
Composition	Single particle	Two or more atoms
Example	O (Oxygen), H (Hydrogen)	O₂ (Oxygen gas), H₂O (Water)

7. What are ions? Differentiate between cations and anions.

Answer:
Ions are charged particles formed when atoms gain or lose electrons.

Types:

Cations

– Positively charged (lose electrons)

Example:

Na⁺, Ca²⁺

Anions

– Negatively charged (gain electrons)

Example:

Cl⁻, SO₄²⁻

8. Define valency. How is it determined?

Answer:
Valency is the combining capacity of an element. It is determined by:

Number of electrons lost or gained

in bonding.

Metals lose electrons

(e.g., Na = 1 valency).

Non-metals gain electrons

(e.g., O = 2 valency).

9. Calculate the molecular mass of water (H₂O).

Answer:
Atomic masses:

Hydrogen (H) = 1 × 2 =

amu

Oxygen (O) =

amu

Total molecular mass =

2 + 16 = 18

amu

10. What is Avogadro’s number? Give its significance.

Answer:
Avogadro’s number = 6.022 × 10²³ particles/mole.

Significance:

Helps in counting

atoms/molecules

Used to calculate

molar volume

Relates

mass, volume, and number of particles

11. Explain the significance of a chemical formula with an example.

Answer:
A chemical formula represents the composition of a compound in terms of:

Elements present in the compound

(e.g., H₂O contains hydrogen and oxygen).

Number of atoms of each element

in a molecule.

Ratio of atoms in a compound

, which is fixed.

Determining molecular mass

by adding atomic masses.

Identifying valency

of elements in the compound.

Understanding chemical bonding

in the molecule.

Example:

The formula of carbon dioxide (CO₂) shows that one carbon atom combines with two oxygen atoms in a

1:2 ratio

12. How do atoms form molecules? Explain different types of molecules with examples.

Answer:
Atoms form molecules by sharing or transferring electrons through chemical bonding.

Types of Molecules:

Monoatomic molecules

– Contain a single atom (e.g., He, Ne).

Diatomic molecules

– Contain two atoms (e.g., O₂, H₂).

Triatomic molecules

– Contain three atoms (e.g., CO₂).

Polyatomic molecules

– Contain more than three atoms (e.g., H₂SO₄, P₄).

Homoatomic molecules

– Made of the same type of atoms (e.g., O₂, N₂).

Heteroatomic molecules

– Made of different types of atoms (e.g., H₂O, NH₃).

Molecules of compounds

– Made of different elements (e.g., NaCl, H₂O).

13. What is the role of valency in the formation of compounds?

Answer:
Valency is the combining capacity of an element and plays a key role in compound formation.

Defines how many bonds an atom can form

(e.g., Na forms one bond, O forms two).

Helps in writing chemical formulas

(e.g., H₂O has

valencies

H = 1, O = 2).

Determines whether an element donates or accepts electrons

(e.g., Na donates, Cl accepts).

Explains why noble gases are unreactive

(valency = 0).

Guides in predicting compound structures

(e.g., CO₂ has linear structure).

Used in balancing chemical equations

(e.g., Na + Cl → NaCl).

Example:

In H₂O,

H (valency 1) combines with O (valency 2) in a 2:1 ratio

14. Differentiate between molecular mass and formula unit mass.

Answer:

Feature	Molecular Mass	Formula Unit Mass
Definition	Sum of atomic masses of all atoms in a molecule	Sum of atomic masses of atoms in an ionic compound
Applies to	Covalent compounds	Ionic compounds
Unit	Atomic mass unit (amu)	Atomic mass unit (amu)
Example	H₂O = 18 amu	NaCl = 58.5 amu
Composition	Can exist independently as a molecule	Exists as a crystal lattice
Type of bonding	Covalent bonds	Ionic bonds
Example compounds	CO₂, H₂O, CH₄	NaCl, K₂SO₄, CaCl₂

15. Explain the concept of a mole with examples.

Answer:
A mole is a unit used to measure the amount of substance, equivalent to 6.022 × 10²³ particles (atoms, molecules, or ions).

1 mole of atoms

6.022 × 10²³ atoms

(e.g., 1 mole of Na has 6.022 × 10²³ Na atoms).

1 mole of molecules

6.022 × 10²³ molecules

(e.g., 1 mole of CO₂ contains 6.022 × 10²³ CO₂ molecules).

Relates atomic/molecular mass to grams

(e.g., 1 mole of O₂ = 32 g).

Helps in chemical calculations

(e.g., moles to mass conversion).

Used in gas laws

(1 mole of any gas = 22.4 L at STP).

Example:

2 moles of CO₂ contain 2 × 6.022 × 10²³ molecules

Application:

Used in

stoichiometry

for balancing reactions.

16. How are atoms of different elements represented? Explain with examples.

Answer:
Atoms of different elements are represented using chemical symbols based on:

First letter of the element

(e.g., H for Hydrogen).

Two letters (if first is common)

(e.g., He for Helium, Hg for Mercury).

Latin names

(e.g., Fe for Iron from

Ferrum

Capital letter for first letter, lowercase for second

(e.g., Na for Sodium).

Uniqueness of symbols helps in writing formulas

Example:

The symbol for carbon is

, and for chlorine is

Application:

Symbols are used in

chemical equations

(e.g., H₂ + O₂ → H₂O).

17. Write the molecular formula and calculate the molecular mass of Ammonia (NH₃).

Answer:
Molecular formula of Ammonia = NH₃

Calculation of molecular mass:

N (Nitrogen)

amu

H (Hydrogen)

amu

× 3 = 3

amu

Total molecular mass = 14 + 3 = 17

amu

18. What are polyatomic ions? Give two examples with their chemical formulas.

Answer:
Polyatomic ions are charged particles containing more than one atom bonded together.

Carry a net charge

(positive or negative).

Act as a single unit in chemical reactions

Can be cations (positive) or anions (negative)

Form ionic compounds

with other ions.

Example:

Sulphate

ion (SO₄²⁻) → Found in

H₂SO₄

Nitrate ion (NO₃⁻) → Found in

NaNO

₃

Application:

Used in acids, salts, and fertilizers.

19. Why do ionic compounds have high melting and boiling points?

Answer:
Ionic compounds have high melting and boiling points due to:

Strong electrostatic forces

between oppositely charged ions.

A large amount of energy

required to break these forces.

Formation of a crystalline structure

, which is highly stable.

High lattice energy

, making them solid at room temperature.

Example:

NaCl melts at

801°C

and boils at

1413°C

Comparison:

Covalent compounds (e.g., water) have lower melting points.

Used in:

Salts, minerals, and industrial processes.

20. Why is water (H₂O) a compound and not a mixture?

Answer:
Water is a compound because:

It has a fixed composition

(H₂:O = 2:1).

Its properties are different

from hydrogen and oxygen.

Formed through a chemical reaction

(H₂ + O₂ → H₂O).

Can’t be separated by physical means

Pure water has a constant boiling point

(100°C).

Follows the law of constant proportions

Example:

Unlike a mixture, water

cannot be separated by filtration or evaporation