Structure of Atom Class 9 Science Revision Notes 2026-27

Structure of Atom Class 9: The chapter on structure of atom class 9 science is one of the most important topics in Physics and Chemistry basics for students. In this chapter, students learn about atoms, subatomic particles, electrons, protons, neutrons, atomic number, mass number, valency, isotopes, and electronic configuration in a simple and easy way. These concepts help students understand how matter is made and why every element behaves differently. The CBSE class 9 syllabus explains the topic step-by-step so students can build a strong base for higher classes and competitive exams also.

These structure of atom class 9 science notes are prepared according to the latest CBSE Board syllabus and exam pattern. The notes include important definitions, formulas, examples, diagrams, and structure of atom class 9 notes questions and answers for quick revision before exams. Students can use these notes for daily study, homework, and last minute preparation. Many learners also search for structure of atom class 9 science notes pdf because short and clear notes make revision faster and less confusing.

In these Class 9 Science CBSE notes, difficult topics are explained in easy words so every student can understand them properly. The chapter may look little confusing at first, but with regular practice and proper revision, it becomes very intresting and scoring for exams.

Introduction to Structure of Atom for Class 9

The structure of the atom has been one of the most investigated questions in the history of science. Understanding it required centuries of thought, experimentation, and revision.

Insight: An atom is the smallest indivisible particle of an element that can take part in a chemical reaction and may or may not exist independently. The word "ATOM" was given by John Dalton.

An element is a pure substance that cannot be broken down into two or more new substances by any chemical or physical means.

Fundamental Particles of an Atom

Every atom is composed of three fundamental (subatomic) particles: electrons, protons, and neutrons.

Comparative Properties

Important Points:

• Electrons reside outside the nucleus and are responsible for chemical bonding. Electrons in the outermost shell are called valence electrons.
• Protons reside in the nucleus and determine the element's identity (atomic number).
• Neutrons reside in the nucleus and contribute to atomic mass but carry no charge.

Proton is a subatomic particle with unit positive charge (+1.602 × 10⁻¹⁹ C) and a mass of 1.6725 × 10⁻²⁷ kg approximately 1837 times greater than the mass of an electron.

In a neutral atom: Number of protons = Number of electrons. Protons and neutrons together are called nucleons.

Thomson's Model of an Atom

Proposed by: J.J. Thomson (1898)

After discovering electrons and protons, Thomson proposed the first model of atomic structure.

The Model:

An atom consists of a sphere of uniform positive electricity, within which electrons are embedded like plums in a pudding (or like seeds in a watermelon). The radius of this sphere is of the order of 10⁻⁸ cm, matching the known size of an atom.

What It Explained:

• The electrical neutrality of atoms: since positive and negative charges are mixed uniformly, they cancel each other out.

Why It Was Discarded:

• Thomson's model could not explain the results of Rutherford's alpha scattering experiment.
• It gave an incorrect picture of how charge is distributed within an atom.

Rutherford's Alpha Scattering Experiment

Scientist: Ernest Rutherford (1911)

This is one of the most famous experiments in the history of atomic physics. It completely overturned the existing understanding of atomic structure.

Experimental Setup

1. Polonium (a radioactive element) was placed inside a lead box with a pinhole. This produced a focused beam of alpha particles (α-particles = He²⁺ nuclei, made up of 2 protons + 2 neutrons).
2. This beam was directed at an extremely thin gold foil (thickness ≈ 0.00004 cm).
3. A circular screen coated with zinc sulphide (ZnS) was placed around the foil. Every time an α-particle struck the screen, it produced a tiny flash of light.

Observations

Conclusions Drawn by Rutherford

The positively charged, heavy, central mass that occupies a very small space within the atom is called the nucleus.

Rutherford's Nuclear Model of Atom

Based on his experiment, Rutherford proposed a new model of atomic structure the planetary or nuclear model.

Main Features:

1. The atom has a small, dense, positively charged nucleus at its centre, which contains almost the entire mass of the atom.
2. Electrons revolve around the nucleus in circular orbits at very high speeds similar to how planets orbit the Sun.
3. The number of electrons in the orbits equals the number of protons in the nucleus.
4. The volume of the nucleus is extremely small compared to the total volume of the atom.
5. Most of the atom is empty space.

Stability Explained: The centrifugal force generated by the fast-moving electrons (acting outward) exactly balances the electrostatic force of attraction pulling them toward the positively charged nucleus maintaining stable orbits.

Defects in Rutherford's Model

Despite its revolutionary significance, Rutherford's model had two serious flaws:

1. Electron distribution not specified: Rutherford did not explain how many electrons could occupy each orbit or what governed their arrangement.
2. Could not explain atomic stability (electromagnetic contradiction): According to classical electromagnetic theory, a charged particle moving in a curved path must continuously radiate energy. As an electron loses energy, it should slow down, spiral inward, and eventually crash into the nucleus causing the atom to collapse. This does not happen in reality, and Rutherford's model offered no explanation.

These limitations were resolved by Neils Bohr's model.

Bohr's Model of an Atom

Proposed by: Neils Bohr (1913)

Bohr accepted Rutherford's nuclear model but added crucial quantisation conditions to explain atomic stability and atomic spectra.

Postulates of Bohr's Theory

1. Fixed Orbits (Energy Levels): Electrons revolve around the nucleus in well-defined circular orbits called shells or energy levels, designated K, L, M, N, ... starting from the one closest to the nucleus.
2. Energy Increases with Distance: The energy associated with electrons in an orbit increases as the distance (radius) of the orbit from the nucleus increases.
3. Quantum Jumps: An electron can jump from a lower energy level to a higher one by absorbing energy, or fall from a higher to a lower level by emitting energy. Electrons do not lose energy while remaining in a fixed orbit.
4. Energy of Radiation: The amount of energy absorbed or emitted during a quantum jump is given by:

ΔE = E₂ − E₁ = hν Where: h = Planck's constant = 6.62 × 10⁻³⁴ Js ν = frequency of the radiation emitted/absorbed E₁ = energy of lower level E₂ = energy of higher level

Why Bohr's model was better: By postulating that electrons in fixed orbits do not radiate energy, Bohr elegantly solved the stability problem that Rutherford's model could not address.

Atomic Structure

An atom has two main structural regions:

(a) Nucleus

• Located at the centre of the atom.
• Contains protons (positively charged) and neutrons (neutral).
• Carries a net positive charge equal to the number of protons.
• Protons + Neutrons = Nucleons (collectively).
• Radius of nucleus ≈ 10⁻¹⁵ m (extremely tiny compared to the atom's radius of ~10⁻¹⁰ m).
• Accounts for nearly all of the atom's mass.

(b) Extra-Nuclear Region

• The region outside the nucleus.
• Contains electrons revolving in fixed energy levels.
• Energy levels are designated K, L, M, N, ... (from innermost to outermost).

Maximum Electrons Per Shell

The maximum number of electrons that can be accommodated in a shell is given by the formula:

Maximum electrons = 2n² (where n = shell number)

Subshells and Orbitals

Each main energy level (shell) is further divided into subshells designated s, p, d, f:

Orbitals are regions within subshells where the probability of finding an electron is maximum. Each orbital holds a maximum of 2 electrons.

Total number of nucleons = Mass number (A) of the atom.

Electronic Configuration

The arrangement of electrons in the different shells of an atom is called its electronic configuration.

Rules for Writing Electronic Configuration

1. Maximum electrons in any shell = 2n².
2. The outermost shell of any atom can hold a maximum of 8 electrons (except the first shell, which holds a maximum of 2).
3. Shells are filled in order, from innermost to outermost.
4. An outermost shell with 8 electrons = Octet (stable).
5. A first shell with 2 electrons = Duplet (stable).

Electronic Configuration of Elements (Atomic Numbers 1–30)

Valency

Valency of an element is the combining capacity of its atoms with atoms of the same or different elements.

Valency arises from the tendency of atoms to achieve a fully filled outermost shell (stable octet or duplet, like noble gases).

How to Determine Valency

The valency depends on the number of valence electrons (electrons in the outermost shell):

If valence electrons = 1 to 4: Valency = Number of valence electrons If valence electrons = 5 to 7: Valency = 8 − Number of valence electrons If valence electrons = 8: Valency = 0 (already stable noble gases)

Examples

Atomic Number (Z)

The atomic number (Z) of an element is the number of protons present in the nucleus of its atom.

Atomic number (Z) = Number of protons
In a neutral atom: Z = Number of protons = Number of electrons
In an ion: Z = Number of protons only (electrons differ)

The atomic number is written as a subscript on the left of the element's symbol:

 ²⁷₁₃Al → Atomic number (Z) = 13, Mass number (A) = 27

Example Sodium (₁₁Na):

• Atomic number = 11
• Number of protons in nucleus = 11
• Charge on nucleus = +11 units
• Number of electrons (neutral atom) = 11

Each element has a unique atomic number it is the fundamental identity of an element. Two different elements can never have the same atomic number.

Mass Number (A)

The mass number (A) of an element is the total number of protons and neutrons in the nucleus.

Mass number (A) = Number of protons (Z) + Number of neutrons (N)
∴ N = A − Z

The mass number is written as a superscript on the left of the element's symbol:

 ²⁷₁₃Al A = 27 (mass number) Z = 13 (atomic number) N = 27 − 13 = 14 neutrons

Relation Between Z, A, and N

A = Z + N N = A − Z Z = A − N

Isotopes

Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A) due to a different number of neutrons.

Same element → Same Z → Same protons → Same electrons → Same chemical properties
Different A → Different neutrons → Different physical properties

Isotopes of Hydrogen

Isotopes of Oxygen

Isotopes of Chlorine

Chlorine has two naturally occurring isotopes: ³⁵₁₇Cl (75%) and ³⁷₁₇Cl (25%).

Average atomic mass of Cl:

= (35 × 75/100) + (37 × 25/100) = 26.25 + 9.25 = 35.5 u

This calculation explains why many elements have fractional atomic masses they are weighted averages of their isotopic masses.

Characteristics of Isotopes

Applications of Radioactive Isotopes

Isobars

Isobars are atoms of different elements with different atomic numbers but the same mass number.

Different Z + Same A → Different elements, different protons/electrons/neutrons, but equal total nucleons

Examples:

Isobars contain different numbers of electrons, protons, and neutrons from each other.

Isotones

Isotones are atoms of different elements that contain the same number of neutrons.

Different elements + Different A and Z, but (A − Z) is the same → Same neutron count

Example:

For ¹³₆C: Neutrons = 13 − 6 = 7
For ¹⁴₇N: Neutrons = 14 − 7 = 7
→ ¹³₆C and ¹⁴₇N are isotones.

Another example: ³⁰₁₄Si, ³¹₁₅P, and ³²₁₆S all have 16 neutrons.

Isoelectronic Species

Atoms, ions, or molecules that have the same number of electrons are called isoelectronic species.

All four species are isoelectronic each has 18 electrons, giving them similar electronic configurations despite being different atoms/ions.

Formulas:

Atomic number (Z) = Number of protons = Number of electrons (neutral atom) Mass number (A) = Number of protons + Number of neutrons = Z + N Number of neutrons = A − Z Electronic configuration rule: Max electrons in shell n = 2n² Valency (valence electrons 1–4) = Valence electrons Valency (valence electrons 5–7) = 8 − Valence electrons Average atomic mass = Σ (mass of isotope × % abundance) / 100 Energy of radiation = ΔE = hν (h = 6.62 × 10⁻³⁴ Js)

Quick Comparison: Isotopes vs Isobars vs Isotones vs Isoelectronic

CBSE Class 9 Science Atoms and Molecules Solved Examples

Example 1 Finding Electrons, Protons, and Neutrons

Problem: Calculate the number of electrons, protons, and neutrons in:
(i) Phosphorus atom (P), (ii) Phosphide ion (P³⁻), (iii) Magnesium ion (Mg²⁺)
(Given: Atomic number of P = 15, mass number = 31; Atomic number of Mg = 12, mass number = 24)

Solution:

(i) Phosphorus atom (neutral):

Electrons = Atomic number = 15 Protons = Atomic number = 15 Neutrons = Mass number − Atomic number = 31 − 15 = 16

(ii) Phosphide ion (P³⁻) gains 3 electrons:

Electrons = 15 + 3 = 18 Protons = 15 (unchanged ions don't gain/lose protons) Neutrons = 31 − 15 = 16 (unchanged)

(iii) Magnesium ion (Mg²⁺) loses 2 electrons:

Electrons = 12 − 2 = 10 Protons = 12 (unchanged) Neutrons = 24 − 12 = 12 (unchanged)

Example 2 Isotope Neutron Count

Problem: An atom has atomic number 41 and mass number 97. How many neutrons are in its isotope of mass number 99?

Solution:

Isotopes have the same atomic number (Z = 41). Neutrons = Mass number − Atomic number = 99 − 41 = 58 neutrons

Example 3 Uranium-238

Problem: For ²³⁸₉₂U, find the number of protons and neutrons.

Solution:

Protons = Atomic number = 92 Neutrons = Mass number − Atomic number = 238 − 92 = 146

Example 4 Finding Atomic Number from Mass Number

Problem: An element has a mass number of 31 and 16 neutrons. Identify the element.

Solution:

Atomic number = Mass number − Neutrons = 31 − 16 = 15 Element with Z = 15 is Phosphorus (P)

Example 5 Average Atomic Mass of Bromine

Problem: Bromine exists as ⁷⁹₃₅Br (49.7%) and ⁸¹₃₅Br (50.3%). Calculate its average atomic mass.

Solution:

Average atomic mass = (79 × 49.7/100) + (81 × 50.3/100) = 39.263 + 40.743 = 80.0 u

Example 6 Electronic Configuration from Shell Data

Problem: An element has 2 electrons in the M-shell. Find its electronic configuration and atomic number.

Solution:

If M-shell has electrons, K and L shells must be completely filled first. K-shell (max) = 2 electrons L-shell (max) = 8 electrons Electronic configuration: K=2, L=8, M=2 → written as 2, 8, 2 Total electrons = 2 + 8 + 2 = 12 Atomic number = 12 → Element is Magnesium (Mg)

Example 7 Determining Valency

Problem: Find the valency of Cl (Z=17), S (Z=16), and Mg (Z=12).

Solution:

Chlorine (Z = 17): Configuration = 2, 8, 7
→ Valence electrons = 7 → Valency = 8 − 7 = 1 (needs 1 electron to complete octet)

Sulphur (Z = 16): Configuration = 2, 8, 6
→ Valence electrons = 6 → Valency = 8 − 6 = 2 (needs 2 electrons to complete octet)

Magnesium (Z = 12): Configuration = 2, 8, 2
→ Valence electrons = 2 → Valency = 2 (loses 2 electrons to achieve stable L-shell octet)