CBSE Notes Class 9 Science Chapter 4 - Structure of the Atom

The diversity of matter comes from different atoms, raising two questions: What differentiates atoms? And, Are atoms divisible or not? The understanding of atomic structure is based on experiments that investigated static electricity and the conditions for electrical conduction in various substances.

Students can also download the ALLEN's App to enhance their learning experiences significantly. It allows students to acquire new information more quickly and efficiently. Click here to understand the advantages of the ALLEN app.

1.0Download CBSE Class 9 Science Chapter 4 Structure of the Atom Notes: Free PDF

Students can now download well-organised and easy-to-understand notes for CBSE Class 9 Science Chapter 4 Structure of the Atom in a free PDF format. These notes are carefully prepared to help students grasp fundamental atomic concepts like subatomic particles, atomic models, isotopes, and isobars with clarity.

2.0Charged Particles in Matter

Rubbing two objects causes them to become electrically charged because atoms contain charged particles. Atoms are composed of three fundamental particles:

• Electron: Negatively charged particles orbiting the nucleus, discovered by J.J. Thomson, who won the Nobel Prize in 1906.
• Proton: Positively charged particles in the nucleus, discovered by Goldstein in 1896.
• Neutron: Neutral particles in the nucleus (except in hydrogen), discovered by James Chadwick in 1932.

3.0The Structure of an Atom

The discovery of electrons and protons disproved Dalton’s claim that atoms were indivisible and indestructible. Subsequently, several scientists proposed atomic models to explain the arrangement of electrons and protons.

Thomson’s Atomic Model (1904):

Known as the "Plum Pudding Model," it aimed to describe atomic structure based on existing knowledge:

1. Electrons are negatively charged.
2. Atoms are neutral overall.

In this model, the atom is envisioned as a sphere of positive charge with negatively charged electrons embedded throughout, similar to plums scattered in a pudding. The electrostatic forces between these particles maintain the atom's neutrality.

Drawbacks of J.J. Thomson’s Atomic Model:

1. It couldn't explain the results of Rutherford's particle scattering experiment.
2. Lacked experimental evidence to support the model.
3. Failed to explain atomic stability.

Rutherford’s Atomic Model (1911)

Rutherford bombarded gold foil with alpha particles (positively charged He²⁺ ions).

Observations:

• Most particles passed straight through.
• Some deflected near the centre.
• A few were reflected.

Conclusions:

• Atoms are mostly empty space, with mass and positive charge concentrated in a tiny nucleus.
• Electrons orbit the nucleus in circular paths, with their movement balanced by centrifugal force.

Nuclear Model of Atom

• The nucleus is the positively charged centre containing most of the atom's mass.
• Electrons orbit the nucleus, which is much smaller than the atom.

Drawbacks:

• The model couldn’t explain atomic stability, as spinning electrons should lose energy and collapse into the nucleus, making matter unstable, which contradicts observations.

Bohr Model of the Atom (1913)

Niels Bohr proposed a model in which electrons orbit the nucleus in quantized energy levels. When they jump to lower-energy orbits, electrons emit radiation. This model explains fixed wavelengths of emitted light and introduces discrete energy levels.

Key Principles:

• Electrons occupy specific energy shells (K, L, M, N, etc.).
• Complete shells make atoms stable.
• Shells are filled in order with a maximum of 2n² electrons (n = shell number).

Electron Distribution:

• K-Shell (n = 1): 2 electrons
• L-Shell (n = 2): 8 electrons
• M-Shell (n = 3): 18 electrons
• N-Shell (n = 4): 32 electrons

Drawbacks:

• Mainly applicable to hydrogen.
• It cannot explain multi-electron atom spectra.
• Ignores wave nature of electrons (de Broglie).
• Does not explain chemical bonding or molecular formation.
• Violates Heisenberg’s Uncertainty Principle.
• It cannot account for the Zeeman and Stark effects.

4.0Neutron

In 1932, J. Chadwick discovered a subatomic particle with no charge and a mass nearly equal to a proton called the neutron. Neutrons, represented by 'n', are found in the nucleus of all atoms except hydrogen. The atomic mass is determined by the sum of the masses of protons and neutrons in the nucleus..

5.0Valency

Valency is determined by the number of electrons in an atom's outermost shell and their need to complete or achieve a stable electron configuration.According to the Bohr-Bury scheme, this shell can hold up to 8 electrons. Atoms with fully-filled outer shells (like helium with 2 electrons or other inert gases with 8) are chemically inactive and have zero valency.

Atoms strive to complete their outer shell (octet) by gaining, losing, or sharing electrons. For example:

• Hydrogen, lithium, and sodium have one outer electron and a valency 1.
• Magnesium has two outer electrons and a valency of 2.
• Aluminum has three outer electrons and a valency of 3.

Atoms close to completing their octet, like fluorine with seven outer electrons, gain 1 electron (valency of 1), while oxygen, with six outer electrons, gains 2 electrons (valency of 2). Valency reflects an atom's combining capacity to achieve a filled outer shell(8 electrons)

6.0Atomic Number  and Mass Number

• The atomic number (Z) is the number of protons in an atom's nucleus and defines the element. For instance, hydrogen has Z = 1, and carbon has Z = 6.
• The mass number is the sum of protons and neutrons in an atom's nucleus. For example, carbon has a mass 12 (6 protons + 6 neutrons), and aluminium has 27 (13 protons + 14 neutrons).

Notation

The atomic number (Z), mass number, and symbol of an element are written in the following format:

7.0Isotopes and Isobars

Isotopes

Some elements have atoms with the same atomic number but different mass numbers, known as isotopes. They are chemically similar but differ in physical properties.

For example, hydrogen has three isotopes: 

If an element has isotopic forms, we calculate its average atomic mass by considering the percentage of each isotope. For chlorine, which consists of 75% ³⁵Cl and 25% ³⁷Cl, the average atomic mass is calculated as:

      =   (75 / 100) × 35u) +(25 / 100 × 37u ) 

      =   35.5 u

This doesn't mean a single chlorine atom has a mass of 35.5 u, but that a sample of chlorine contains both isotopes and the weighted average mass is 35.5 u.

Isobars

Atoms of different elements with distinct atomic numbers but the same mass number are called isobars. For example, calcium (atomic number 20) and argon (atomic number 18) have different electrons, yet both have a mass 40. This means the number of nucleons (protons and neutrons) is the same in both atoms.

8.0Key Features of CBSE Science Notes for Class 9 Chapter 4 Structure of the Atom

• Complete Chapter Coverage: Includes all important topics like atomic models, subatomic particles, electronic configuration, valency, and isotopes.
• Concept-Based Learning: Each concept is explained with real-life examples, diagrams, and simplified language to ensure better understanding.
• Updated as per the Latest CBSE Syllabus: Notes strictly follow the current NCERT Solutions guidelines and CBSE curriculum.
• Easy-to-Understand Format: Well-organized notes with bullet points, flowcharts, and illustrations make learning fun and effective.