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 E=hνE = 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 4 prefer to share 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:

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

Cation (+ve ion) is formed when an atom loses electrons (e.g., Na → Na⁺ + e⁻).

Anion (-ve 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, Ar 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.