Extra 30 short questions and answers from the chapter:3 "Atoms and Molecules" for Class 9 CBSE Science:
1. Define the Law of Conservation of Mass with an example.
Answer:
The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction.
Example:
When hydrogen (H₂) reacts with oxygen (O₂) to form water (H₂O), the total mass of reactants and products remains the same:
2H2+O2→2H2O2H₂ + O₂ → 2H₂O2H2+O2→2H2O
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 the same proportion by mass.
Example:
Water (H₂O) always contains hydrogen and oxygen in a fixed ratio of 1:8 by mass, regardless of the source.
3. What are atoms? Name two scientists who contributed to atomic theory.
Answer:
Atoms are the smallest particles of an element that cannot be divided further by chemical means.
Scientists:
John Dalton
– Proposed the
Atomic Theory
.
Ernest Rutherford
– Discovered the
nucleus
in an atom.
4. Define atomic mass unit (amu) and give two examples.
Answer:
1 atomic mass unit (amu) is 1/12th the mass of a carbon-12 atom.
Examples:
Hydrogen (H) =
1
amu
Oxygen (O) =
16
amu
5. What is a molecule? Give an example of a monoatomic and diatomic molecule.
Answer:
A molecule is the smallest unit of a substance that retains its chemical properties.
Examples:
Monoatomic molecule:
He (Helium)
Diatomic molecule:
O₂ (Oxygen gas)
6. Differentiate between an element and a compound.
Answer:
|
Feature |
Element |
Compound |
|---|---|---|
|
Definition |
A pure substance made of one kind of atom |
A substance formed by chemically combining two or more elements |
|
Composition |
Cannot be broken down further |
Can be broken down into elements |
|
Examples |
Oxygen (O₂), Iron (Fe) |
Water (H₂O), Carbon dioxide (CO₂) |
7. What is a chemical formula? Write the chemical formula for water and ammonia.
Answer:
A chemical formula represents the composition of a compound using symbols and numbers.
Examples:
Water =
H₂O
Ammonia =
NH₃
8. What is an ion? Give an example of a cation and an anion.
Answer:
An ion is a charged particle formed when an atom gains or loses electrons.
Examples:
Cation (Positive ion):
Na⁺ (Sodium ion)
Anion (Negative ion):
Cl⁻ (Chloride ion)
9. What is valency? How is it determined?
Answer:
Valency is the combining capacity of an atom based on the number of electrons it gains, loses, or shares.
Determined by:
Metals:
Lose electrons (
e.g., Na = 1 valency
)
Non-metals:
Gain electrons (
e.g., Oxygen = 2 valency
)
10. What is Avogadro’s number? What does it represent?
Answer:
Avogadro’s number is 6.022 × 10²³ particles per mole. It represents:
Atoms in 1 mole of an element
Molecules in 1 mole of a compound
11. Define molecular mass with an example.
Answer:
Molecular mass is the sum of atomic masses of all atoms in a molecule.
Example: CO₂
= (1 × 12) + (2 × 16)
= 44 amu
12. What is a mole? Give an example.
Answer:
A mole is the amount of a substance containing 6.022 × 10²³ particles.
Example: 1 mole of water (H₂O) contains 6.022 × 10²³ molecules.
13. Write the molecular formula of the following compounds: (a) Sulfuric acid (b) Glucose
Answer:
Sulfuric acid
=
H₂SO₄
Glucose
=
C₆H₁₂O₆
14. How is molecular mass different from atomic mass?
Answer:
|
Feature |
Atomic Mass |
Molecular Mass |
|---|---|---|
|
Definition |
Mass of a single atom |
Sum of atomic masses in a molecule |
|
Unit |
amu |
amu |
|
Example |
Oxygen = 16 amu |
CO₂ = 44 amu |
15. What is the molecular mass of methane (CH₄)?
Answer:
Atomic masses:
Carbon (C) = 12
Hydrogen (H) = 1 × 4 = 4
Total molecular mass =
12 + 4 = 16
amu
16. Why do noble gases have zero valency?
Answer:
Noble gases have a completely filled outer shell, so they do not gain or lose electrons, making their valency zero.
Example: Helium (He), Neon (Ne).
17. Why is water called a compound and not a mixture?
Answer:
Water (H₂O) is a compound because:
It has
fixed composition
(H₂:O = 2:1).
Its properties
differ from its elements (H and O)
.
It cannot be separated by
physical means
.
18. What are polyatomic ions? Give two examples.
Answer:
Polyatomic ions are ions containing more than one atom.
Examples:
Sulphate
ion (SO₄²⁻)
Nitrate ion (NO₃⁻)
19. Differentiate between empirical and molecular formulas.
Answer:
|
Feature |
Empirical Formula |
Molecular Formula |
|---|---|---|
|
Definition |
Simplest ratio of atoms |
Actual number of atoms |
|
Example |
CH₂ (for C₂H₄) |
C₂H₄ |
20. What is gram molecular mass?
Answer:
The molecular mass of a substance expressed in grams.
Example:
Molecular mass of CO₂ = 44 amu
Gram molecular mass = 44 g
21. What is the significance of the formula of a compound?
Answer:
The chemical formula of a compound provides:
The
types of elements
present in the compound.
The
number of atoms
of each element.
The
ratio
in which the atoms combine.
It helps in determining the
molecular mass
.
Example: In H₂O, two hydrogen (H) atoms and one oxygen (O) atom are present in a fixed ratio of 2:1.
22. 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) |
23. What is the formula unit mass? How is it calculated?
Answer:
Formula unit mass is the
sum of atomic masses
of all atoms in a compound’s formula unit.
It is used for
ionic compounds
.
Example: NaCl
= Na (23) + Cl (35.5)
= 58.5 amu
24. Why is the atomic mass of chlorine taken as 35.5 u and not a whole number?
Answer:
The atomic mass of chlorine is 35.5 u because chlorine has two isotopes:
Cl-35 (75%)
Cl-37 (25%)
The weighted average of these isotopes gives
35.5 u
.
25. How do you calculate the number of moles in a given mass?
Answer:
The number of moles (nnn) is calculated using:
n=Given mass (g)Molar mass (g/mol)n = \frac{\text{Given mass (g)}}{\text{Molar mass (g/mol)}}n=Molar mass (g/mol)Given mass (g)
Example:
For 88 g of CO₂,
Molar mass of CO₂ = 44 g/mol
n=8844=2 moles n = \frac{88}{44} = 2 \text{ moles}n=4488=2 moles
26. What is the molecular mass of calcium carbonate (CaCO₃)?
Answer:
Molecular mass = sum of atomic masses of all atoms in the molecule.
Ca = 40
C = 12
O₃ = 16 × 3 = 48
Total =
40 + 12 + 48 = 100
amu
27. Define molar mass and give two examples.
Answer:
Molar mass is the mass of one mole of a substance in grams per mole (g/mol).
Examples:
H₂O =
18 g/mol
CO₂ =
44 g/mol
28. Find the mass of 2 moles of oxygen gas (O₂).
Answer:
Molar mass of O₂ = 32 g/mol
Mass = Number of moles × Molar mass
=2×32=64g= 2 × 32 = 64 g=2×32=64g
29. Differentiate between molecular mass and molar mass.
Answer:
|
Feature |
Molecular Mass |
Molar Mass |
|---|---|---|
|
Definition |
Mass of one molecule |
Mass of one mole of a substance |
|
Unit |
amu |
g/mol |
|
Example |
CO₂ = 44 amu |
CO₂ = 44 g/mol |
30. How many molecules are present in 5 moles of water (H₂O)?
Answer:
Using Avogadro’s number:
