Class 9 Chapter 4 Long Qa Science Cbse Ncert

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Extra 20 important long-answer questions from the "Structure of the Atom" chapter:4 of Class 9 CBSE Science:-

1. Explain Dalton’s Atomic Theory and its limitations.

Answer: Dalton proposed the first atomic theory:

All matter is made of tiny, indivisible particles called atoms.

Atoms of the same element are identical in mass and properties.

Atoms of different elements differ in mass and properties.

Atoms combine in simple whole-number ratios to form compounds.

Chemical reactions involve the rearrangement of atoms.

Limitations:

Atoms are divisible (discovery of protons, neutrons, and electrons).

Isotopes exist (same element, different mass numbers).

Atoms can change in nuclear reactions.

2. Describe Rutherford’s gold foil experiment and its conclusions.

Answer:

Rutherford bombarded thin gold foil with alpha (α) particles.

Most particles passed through, some deflected, and few bounced back.

This led to the discovery of the

nucleus

in an atom.

Conclusions:

Atom has a

small, dense, positively charged nucleus

Electrons revolve around the nucleus in empty space.

Most of the atom is empty space.

Limitations:

Could not explain electron stability.

Did not describe the arrangement of electrons.

3. State Bohr’s atomic model and its advantages.

Answer:

Electrons revolve in

fixed energy levels (shells)

around the nucleus.

Electrons do not lose energy while revolving.

Energy is absorbed or emitted when electrons jump between levels.

The energy of an orbit is

quantized

The

maximum number of electrons in a shell

follows the

2n²

rule.

Advantages:

Explained atomic stability.

Justified hydrogen’s spectral lines.

Introduced the concept of quantized energy levels.

4. Differentiate between Isotopes and Isobars with examples.

Answer:

Property	Isotopes	Isobars
Definition	Atoms of the same element with different mass numbers	Atoms of different elements with the same mass number
Atomic Number	Same	Different
Mass Number	Different	Same
Examples	Carbon-12, Carbon-14	Argon-40, Calcium-40

5. Define valency. How is it determined? Give examples.

Answer:

Valency

is the combining capacity of an element.

It is determined by the number of electrons

lost, gained, or shared

If an element

loses electrons

, valency = number of lost electrons.

If it

gains electrons

, valency = number of gained electrons.

Example: Na (2,8,1) loses 1 electron, so valency = 1.

Cl (2,8,7) gains 1 electron, so valency = 1.

6. Explain the discovery of electrons, protons, and neutrons.

Answer:

Electron (J.J. Thomson, 1897)

: Cathode ray experiment.

Proton (Goldstein, 1886)

: Canal ray experiment.

Neutron (James Chadwick, 1932)

: Bombarded beryllium with alpha particles.

7. What is the electronic configuration? Explain with an example.

Answer:

Electronic configuration

is the arrangement of electrons in shells.

It follows the

2n²

rule.

Example: Sodium (Z=11) →

2,8,1

The last shell determines valency.

Noble gases

have a complete outer shell (stable configuration).

8. Compare Rutherford’s and Bohr’s models of the atom.

Answer:

Feature	Rutherford’s Model	Bohr’s Model
Electron Path	Random motion	Fixed orbits
Energy Loss	Predicted loss	No energy loss
Stability	Could not explain	Explained
Nucleus	Present	Present
Energy Levels	Not defined	Quantized

9. Explain the concept of atomic number and mass number.

Answer:

Atomic number (Z)

= Number of protons = Number of electrons.

Mass number (A)

= Protons + Neutrons.

Example: Oxygen (Z=8, A=16) has 8 protons and 8 neutrons.

Atomic number determines the element’s identity.

10. Describe the structure of an atom with a labelled diagram.

Answer:

Atom has

a central nucleus

containing protons and neutrons.

Electrons

revolve around the nucleus

in shells.

The

first shell (K) holds 2 electrons, L holds 8, M holds 18

, etc.

The

nucleus is positively charged

, and the atom is neutral.

11. Explain J.J. Thomson’s Model of an Atom. What were its limitations?

Answer:

J.J. Thomson proposed the

Plum Pudding Model

in 1897.

Atoms consist of a positively charged sphere with negatively charged electrons embedded in it.

The positive charge balances the negative charge, making the atom neutral.

Electrons were thought to be like "plums" in a positively charged "pudding."

The model explained the neutrality of atoms.

Limitations:
6. It could not explain the presence of a nucleus.
7. It failed to explain how electrons were arranged in an atom.

12. Describe the properties of cathode rays and how their discovery led to the identification of electrons.

Answer:

Cathode rays

were discovered by

J.J. Thomson

using a discharge tube experiment.

Cathode rays originate from the

negative electrode (cathode)

in a vacuum tube.

They travel in

straight lines

and cast shadows.

They are

negatively charged

, as they were deflected towards a positive plate.

They produce a

mechanical effect

, rotating a light paddle wheel.

This led to the discovery of the

electron

, a negatively charged particle.

The charge-to-mass ratio (e/me/me/m) of an electron was calculated.

13. Explain the significance of neutrons in an atom.

Answer:

Neutrons were discovered by

James Chadwick in 1932

Neutrons have

no charge

(neutral) and are present in the nucleus.

They provide

stability

by reducing repulsion between protons.

The number of neutrons affects the

mass number

of an atom.

Elements with different numbers of neutrons are called

isotopes

Neutrons play a key role in

nuclear reactions

(e.g., fission and fusion).

Atoms with an unstable neutron-to-proton ratio become

radioactive

14. What are the limitations of Rutherford’s atomic model?

Answer:

Rutherford proposed that electrons move in

random orbits

around the nucleus.

According to classical physics, moving electrons should lose energy and

spiral into the nucleus

This would make atoms

unstable

, but atoms are actually stable.

It did not explain the arrangement of

electrons in energy levels

It could not explain the

line spectra of elements

It failed to address why electrons do not

fall into the nucleus

Bohr’s model later resolved these issues by introducing

energy levels

15. Explain the relationship between energy levels and the spectrum of an atom.

Answer:

Electrons occupy

specific energy levels

in an atom.

When an electron

absorbs energy

, it jumps to a

higher energy level

(excited state).

When it returns to a

lower energy level

, it

releases energy

as light.

The emitted light forms

spectral lines

, which are unique for each element.

This explains the

line spectra

observed in atomic emission spectra.

The energy of the emitted photons follows

Bohr’s equation

hνE

hν

The spectra help identify

elements in stars and chemical samples

16. Explain the concept of quantum numbers and their significance.

Answer:

Quantum numbers describe the

position and energy

of electrons in an atom.

Principal quantum number (n)

: Indicates the

energy level

(K, L, M, N).

Azimuthal quantum number (l)

: Determines the

shape of the orbital

(s, p, d, f).

Magnetic quantum number (m)

: Describes the

orientation of the orbital

Spin quantum number (s)

: Determines the

spin direction

of the electron (+½ or -½).

They help understand

electronic configuration

and

chemical bonding

Quantum numbers explain the

periodic trends

in the periodic table.

17. How does the concept of valency help in chemical bonding?

Answer:

Valency

is the number of electrons an atom

gains, loses, or shares

Atoms with

1, 2, or 3 valence electrons

tend to

lose

them, forming

positive ions

Atoms with

5, 6, or 7 valence electrons

tend to

gain

electrons, forming

negative ions

Elements with a valency of

prefer to

electrons (e.g., carbon in methane).

Noble gases have a

valency of 0

because they have a

full outer shell

Valency determines the

types of bonds

(ionic, covalent).

It helps predict

chemical reactivity

and

compound formation

18. How do isotopes of an element differ in physical and chemical properties?

Answer:

Physical properties differ

because isotopes have different

mass numbers

Chemical properties remain similar

because they have the same

electronic configuration

Isotopes may have different

densities and boiling points

Radioactive isotopes

are unstable and decay over time (e.g., Carbon-14).

Applications

: Uranium-235 is used in nuclear reactors, while Carbon-14 is used for dating fossils.

Some isotopes are used in

medical treatments

, e.g., Cobalt-60 for cancer therapy.

Example

: Hydrogen has

three isotopes

– Protium (¹H), Deuterium (²H), and Tritium (³H).

19. Describe the formation of ions with examples.

Answer:

ion

is a charged particle formed when an atom gains or loses electrons.

Cation

ion) is formed when an atom

loses electrons

(e.g., Na → Na⁺ + e⁻).

Anion

ion) is formed when an

atom

gains electrons

(e.g., Cl + e⁻ → Cl⁻).

Metals usually form

cations

, while non-metals form

anions

Ionic bonds

are formed when cations and anions attract each other (e.g., NaCl).

Examples

Mg²⁺ forms when magnesium loses 2 electrons.

O²⁻ forms when oxygen gains 2 electrons.

Ions play an important role in

electrolysis and conductivity

20. Explain the stability of atoms based on electronic configuration.

Answer:

Atoms are stable when they have a

full outermost shell (octet rule)

Noble gases like

He, Ne,

are naturally stable due to their full shells.

Atoms with

incomplete outer shells

are unstable and react to form

bonds

Metals

lose electrons

to attain a noble gas configuration.

Non-metals

gain electrons

to complete their octet.

Carbon forms

covalent bonds

as it has four valence electrons.

The concept of stability explains

why elements form compounds