CBSE Class 9 Science Revision Notes Chapter 8: Journey Inside the Atom

These Journey Inside the Atom notes explain the development of atomic models, the nucleus, energy levels and the arrangement of electrons. The chapter also covers electrons, protons and neutrons, atomic number and mass number, electronic configuration, valency, isotopes and isobars.

Everything around us is made of matter, and matter is made of atoms. Scientists once believed that atoms were indivisible, but later experiments showed that an atom contains smaller particles and has an organised internal structure.

These CBSE Class 9 Science Chapter 8 notes explain how the idea of the atom developed from Dalton’s model to the Thomson, Rutherford and Bohr models. The chapter also covers subatomic particles, electronic configuration, valency, average atomic mass and the difference between isotopes and isobars.

Key Takeaways

• Nucleus: It is the small, dense centre of the atom containing protons and neutrons.
• Atomic number: It is equal to the number of protons in an atom.
• Mass number: It is the total number of protons and neutrons.
• Valency: It depends on the electrons lost, gained or shared to achieve stability.

Access Class 9 Science Chapter 8 Journey Inside the Atom Notes in 30 Minutes

Revise the chapter in three parts:

• First 10 minutes: Dalton, Thomson, Rutherford and Bohr atomic models
• Next 10 minutes: Electrons, protons, neutrons, atomic number, mass number and electronic configuration
• Final 10 minutes: Valence electrons, valency, isotopes, average atomic mass and isobars

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Structure of Atom Class 9: Development of Atomic Theory

The structure of atom Class 9 chapter traces how the scientific understanding of atoms changed over time.

Early thinkers suggested that matter could be divided only up to a certain point. Later, experiments helped scientists build and improve different atomic models.

Acharya Kanada and Parmanu

Acharya Kanada proposed that repeated division of matter would finally produce extremely small indivisible particles.

He called these particles parmanus.

According to his idea, parmanus combined to form larger units and all material substances. However, this explanation was philosophical rather than experimental.

Greek Concept of Atomos

Greek philosophers Leucippus and Democritus proposed a similar idea.

They used the word atomos, which means indivisible.

These early ideas were important, but they were not based on scientific experiments.

Dalton’s Atomic Theory

In 1808, John Dalton presented the first scientific theory of atoms.

According to Dalton:

• All matter is made of atoms.
• Atoms were considered indivisible.
• Atoms of the same element were believed to be identical.
• Atoms of different elements were different.
• Atoms combined in fixed ratios to form compounds.

Dalton’s theory gave a scientific base to the study of matter.

Later discoveries showed that atoms contain smaller particles, so the idea of an indivisible atom had to be modified.

Thomson Atomic Model and Discovery of Electrons

In 1897, J. J. Thomson studied how electric current passed through gases at low pressure.

He used a cathode ray tube containing two electrodes. He observed rays moving from the negative electrode, called the cathode, towards the positive electrode, called the anode.

Thomson concluded that these rays contained negatively charged particles.

These particles were later called electrons.

Important Findings About Electrons

• Electrons carry negative charge.
• Their mass is much smaller than the mass of an atom.
• They are present in atoms of every element.
• The relative charge of an electron is taken as −1.

The fact that cathode rays did not depend on the gas or electrode material showed that electrons were a common part of all atoms.

Thomson Atomic Model

The Thomson atomic model described an atom as a sphere of positive charge with negatively charged electrons distributed throughout it.

This model is often called the plum pudding model.

It can also be compared to a watermelon:

• The red pulp represents the positive charge.
• The seeds represent electrons.
• The positive and negative charges balance each other.

This model explained why an atom is electrically neutral.

Limitation of Thomson’s Model

Thomson’s model could not explain the observations of the Rutherford gold foil experiment.

It treated the positive charge as uniformly spread throughout the atom, but later evidence showed that the positive charge was concentrated in a small region.

Rutherford Gold Foil Experiment and Atomic Model

In 1911, Geiger and Marsden performed the Rutherford gold foil experiment under Ernest Rutherford’s guidance.

They directed a narrow beam of positively charged alpha particles towards a very thin sheet of gold.

According to Thomson’s model, most alpha particles should have passed through with only small deflections.

Observations of the Gold Foil Experiment

• Most alpha particles passed straight through.
• Some were deflected through small angles.
• Very few were deflected through large angles.
• A tiny number bounced back.

The change in the straight path of alpha particles is called scattering.

Conclusions from Rutherford’s Experiment

Rutherford concluded:

• Most of the atom is empty space.
• Positive charge is concentrated in a very small central region.
• This central region is called the nucleus.
• Most of the atom’s mass is present in the nucleus.
• Electrons move around the nucleus.

The diameter of an atom is about 10⁻¹⁰ m, while the nucleus is about 10⁻¹⁵ m.

This means that the nucleus is nearly one lakh times smaller than the atom.

Rutherford Atomic Model and Its Limitation

The Rutherford atomic model is also known as the planetary model.

It proposed that electrons move around a small, positively charged nucleus, similar to planets revolving around the Sun.

Main Features of Rutherford’s Model

• The nucleus is present at the centre.
• It contains all the positive charge.
• Most of the atom’s mass lies in the nucleus.
• Electrons revolve around the nucleus.
• Most of the atom is empty space.

Limitation of Rutherford’s Atomic Model

An electron moving in a circular path is constantly changing direction. This means that it is accelerating.

According to classical ideas, an accelerating charged particle should lose energy.

If the electron kept losing energy, it would move in a spiral path and eventually fall into the nucleus.

This would make the atom unstable.

Since atoms are stable, Rutherford’s model could not fully explain atomic structure.

Discovery of Protons

Rutherford showed that the positive charge of the nucleus comes from particles called protons.

A proton:

• Has a relative charge of +1
• Is present inside the nucleus
• Has much greater mass than an electron

In a neutral atom, the number of protons equals the number of electrons.

Examples:

• Helium has 2 protons and 2 electrons.
• Sodium has 11 protons and 11 electrons.

Their positive and negative charges balance each other.

Bohr Atomic Model and Electron Shells

In 1913, Niels Bohr proposed a new model to explain atomic stability.

The Bohr atomic model introduced fixed paths for electrons.

These fixed paths are called shells, orbits or energy levels.

Main Features of Bohr’s Model

• Electrons move in fixed circular paths around the nucleus.
• Each shell has a fixed amount of energy.
• Electrons do not lose energy while moving in an allowed shell.
• Electrons cannot remain between two shells.
• An electron can move to another shell by absorbing or releasing a fixed amount of energy.

Names of Electron Shells

The shells are represented as:

• K shell: n = 1
• L shell: n = 2
• M shell: n = 3
• N shell: n = 4

The K shell is nearest to the nucleus and has the lowest energy.

The energy of the shells increases with distance from the nucleus.

Bohr’s model explained atomic stability better than Rutherford’s model.

Electrons, Protons and Neutrons in an Atom

Electrons, protons and neutrons are the three main subatomic particles discussed in this chapter.

The phrase electrons protons and neutrons is commonly searched by students, but in correct scientific writing the terms are written with commas.

Discovery of the Neutron

James Chadwick discovered the neutron in 1932.

A neutron:

• Has no electrical charge
• Has mass nearly equal to that of a proton
• Is found in the nucleus
• Is absent in the common hydrogen atom

The discovery of neutrons helped explain why atomic mass could not be calculated using protons alone.

Contribution to Atomic Mass

Protons and neutrons together are called nucleons.

Most of an atom’s mass comes from nucleons.

The mass of an electron is very small and is generally ignored in basic calculations.

Symbols of Elements in Class 9 Science Chapter 8

Chemical symbols provide a common and standard way to represent elements.

The International Union of Pure and Applied Chemistry approves the names and symbols of elements.

Rules for Writing Chemical Symbols

• A symbol may contain one or two letters.
• The first letter is always uppercase.
• The second letter, if present, is lowercase.
• Some symbols come from English names.
• Some come from Latin, Greek or German names.

These symbols are internationally recognised and allow scientists to communicate clearly.

Atomic Number and Mass Number

The concepts of atomic number and mass number help describe the composition of an atom.

Atomic Number

The number of protons present in the nucleus is called the atomic number.

It is represented by Z.

Atomic number = Number of protons

In a neutral atom:

Number of protons = Number of electrons

The atomic number determines the identity of an element.

Examples:

• Hydrogen: Z = 1
• Helium: Z = 2
• Lithium: Z = 3
• Carbon: Z = 6

Mass Number

The total number of protons and neutrons present in the nucleus is called the mass number.

It is represented by A.

Mass number = Number of protons + Number of neutrons

A = p + n

Therefore:

Number of neutrons = Mass number − Atomic number

n = A − Z

Standard Atomic Notation

An atom is represented as:

ᴬZX

Where:

• X = symbol of the element
• A = mass number
• Z = atomic number

For carbon:

¹²₆C

This means:

• Atomic number = 6
• Protons = 6
• Electrons = 6
• Mass number = 12
• Neutrons = 12 − 6 = 6

Numerical Example

An atom has:

• Atomic number = 26
• Mass number = 56

Number of protons = 26

Number of electrons = 26

Number of neutrons = 56 − 26

Number of neutrons = 30

Electronic Configuration Class 9 and Bohr-Bury Rules

The electronic configuration Class 9 topic explains how electrons are distributed in different shells.

The Bohr-Bury rules are used to place electrons in the K, L, M and N shells.

Rule 1: Maximum Number of Electrons

The maximum number of electrons in a shell is given by:

Maximum number of electrons = 2n²

Where n is the shell number.

Rule 2: Outermost Shell

The outermost shell can contain a maximum of 8 electrons.

The first shell can contain only 2 electrons.

Rule 3: Filling Order

Electrons fill the inner shells before moving to the outer shells.

The filling order is:

K → L → M → N

The L shell begins filling only after the K shell is complete.

Electronic Configuration of the First 18 Elements

Valency and Valence Electrons

The concepts of valency and valence electrons explain how atoms combine with one another.

The outermost shell containing electrons is called the valence shell.

The electrons present in it are called valence electrons.

Stable Electronic Configuration

Atoms with a completely filled outermost shell are generally stable.

• Helium is stable with 2 electrons.
• Neon and argon are stable with 8 electrons.

A valence shell containing 8 electrons is called an octet.

Meaning of Valency

Valency is the combining capacity of an atom.

It is equal to the number of electrons lost, gained or shared to achieve a stable outer shell.

Elements with One to Four Valence Electrons

Elements with fewer than four valence electrons generally lose electrons.

Examples:

• Sodium: 2, 8, 1 → valency 1
• Magnesium: 2, 8, 2 → valency 2
• Aluminium: 2, 8, 3 → valency 3

Elements with Five to Seven Valence Electrons

These atoms generally gain electrons.

Their valency can be calculated as:

Valency = 8 − Number of valence electrons

Examples:

• Nitrogen: 2, 5 → valency 3
• Oxygen: 2, 6 → valency 2
• Fluorine: 2, 7 → valency 1

Valency of Carbon

Carbon has the electronic configuration 2, 4.

It cannot easily gain or lose four electrons. It usually shares four electrons.

Therefore, its valency is 4.

Isotopes and Isobars

The chapter explains isotopes and isobars as two different relationships between atoms.

Isotopes

Isotopes are atoms of the same element that have:

• The same atomic number
• The same number of protons
• Different numbers of neutrons
• Different mass numbers

Since their electronic configurations are the same, isotopes generally have similar chemical properties.

Their physical properties may differ.

Isotopes of Hydrogen

All three isotopes have one electron in a neutral atom.

Isotopes of Carbon

Carbon has three isotopes:

• ¹²₆C
• ¹³₆C
• ¹⁴₆C

Each isotope has 6 protons and 6 electrons.

They differ in the number of neutrons.

Uses of Isotopes

Average Atomic Mass of an Element

The average atomic mass of an element depends on the masses and relative abundances of its naturally occurring isotopes.

Example of Chlorine

Chlorine occurs mainly as:

• Chlorine-35: about 75%
• Chlorine-37: about 25%

Average atomic mass:

Average atomic mass = (35 × 75/100) + (37 × 25/100)

= 26.25 + 9.25

Average atomic mass = 35.5 u

This does not mean that one chlorine atom has a mass of 35.5 u.

It represents the weighted average mass of a large number of naturally occurring chlorine atoms.

Isobars in Class 9 Science Chapter 8

Isobars are atoms of different elements that have:

• The same mass number
• Different atomic numbers
• Different numbers of protons

Examples include:

• Argon-40: atomic number 18
• Potassium-40: atomic number 19
• Calcium-40: atomic number 20

All have the same mass number but belong to different elements.

Difference Between Isotopes and Isobars

Comparison of Atomic Models

Class 9 Atomic Structure Formulas

The following Class 9 atomic structure formulas and relations are important for revision:

Journey Inside the Atom Notes: Quick Revision

These Journey Inside the Atom notes summarise the major ideas of the chapter:

• Atoms are the building blocks of matter.
• Dalton treated atoms as indivisible particles.
• Thomson discovered electrons and proposed the plum pudding model.
• Rutherford’s experiment led to the discovery of the nucleus.
• Most of the atom is empty space.
• Bohr proposed fixed energy levels.
• Electrons are present outside the nucleus.
• Protons and neutrons are found inside the nucleus.
• Atomic number equals the number of protons.
• Mass number equals protons plus neutrons.
• Electronic configuration follows the Bohr-Bury rules.
• Valency depends on electrons gained, lost or shared.
• Isotopes have the same atomic number but different mass numbers.
• Isobars have the same mass number but different atomic numbers.
• Average atomic mass depends on the abundance of isotopes.

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