⚛ Chapter 3 · NCERT Science IX

ATOMS &
MOLECULES

From Dalton's indivisible spheres to Avogadro's number — the bedrock of all chemistry, one atom at a time.

n = m/M N = n × 6.022×10²³ 22.4 L/mol
3Key Laws
6.022×10²³Avogadro
★★★★★Exam Weight
📍 Chapter Snapshot
📖Chapter 03
⏱️Est. Study Time: 6 hrs
🎯Difficulty: High
📝Numericals: Yes — Many
Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical reaction. (Lavoisier, 1774)
Law of Constant Proportions
A pure compound always has elements in a fixed mass ratio. (Proust, 1799)
Dalton's Atomic Theory
Atoms are indivisible, indestructible building blocks. Each element has unique atoms.
Avogadro's Law
Equal volumes of gases at same T & P contain equal number of molecules.
🎯 Why This Chapter Matters for Exams
📌 Mole concept numericals appear in every board exam — master the formula triangle.
📌 Writing chemical formulae of compounds (with valency) is a guaranteed 2-mark question.
📌 Atomic mass vs molecular mass distinctions are classic MCQ fodder.
📌 Laws of chemical combination carry full 5-mark theory questions in SA-I/SA-II.
💡 Key Concept Highlights
AtomMoleculeMoleAvogadro NumberAtomic MassMolecular MassMolar MassValencyIonPolyatomic IonChemical FormulaAMU (u)
⚗️ Important Formula Capsules
Moles of Substancen = m / Mmol
Number of ParticlesN = n × Nₐparticles
Avogadro NumberNₐ = 6.022 × 10²³mol⁻¹
Molar Volume (gas)22.4 L at STPL/mol
📚 What You Will Learn
  • 1Laws of chemical combination and their historical significance
  • 2Writing and balancing chemical formulae using valency
  • 3Calculating atomic, molecular, and formula masses
  • 4Mole concept and its applications in stoichiometry
  • 5Avogadro's number and molar volume at STP
🔗 Navigate to Chapter Resources
📄
Full Notes
🧩
MCQs Practice
✅
True / False
📝
NCERT Exercises
🏆 Exam Strategy & Preparation Tips
01
Mole Triangle
Draw a triangle: Mass, Moles, Particles. Cover any corner to get the formula.
02
Valency Table
Memorise valencies of 20 common ions — it unlocks all formula-writing questions.
03
Step-by-Step
Show all working in numericals. Partial marks are given even for wrong final answers.
04
Unit Vigilance
Always write units: g/mol, u, mol — unit errors cost marks in board exams.
Chapter 3 · CBSE · Class IX
⚛️

Kanad's Atomic Theory

NCERT Class 9 Science Atoms Molecules Laws of Chemical Combination Law of Conservation of Mass Law of Definite Proportions Dalton Atomic Theory Atomic Structure Molecules of Elements Molecules of Compounds Ions Cations Anions Chemical Formula Molecular Formula Writing Chemical Formulae Valency Mole Concept Avogadro Number Molar Mass Atomic Mass Unit Relative Atomic Mass Molecular Mass Formula Mass Mass of a Mole Balancing Chemical Equations Symbols of Elements Polyatomic Ions Combination Reaction Decomposition Reaction
📋
Table of Contents
15 sections
🗺️ Overview
Kanad's Atomic Theory – The Earliest Concept of Atoms
Long before the development of modern chemistry, ancient philosophers attempted to understand the nature of matter. Among them, Maharishi Kanad, an Indian philosopher and founder of the Vaisheshika School of Philosophy, proposed one of the earliest known atomic theories around 500 BCE. His ideas laid the foundation for the concept that all matter is composed of extremely small particles called Parmanu (atoms).
🖼️ Figure
Maharishi Kanad Founder of Early Atomic Theory
Maharishi Kanad (approximately 500 BCE), one of the earliest thinkers to propose an atomic model of matter.
🏛️ Historical Note
Historical Background

During ancient times, philosophers observed that substances could be broken into smaller pieces. Kanad questioned whether this process could continue indefinitely. He concluded that a stage must eventually be reached where further division is impossible. The smallest indivisible particle obtained was called a Parmanu.

This idea emerged nearly two thousand years before modern scientists such as John Dalton proposed scientific atomic theory based on experimental evidence.

📘 Definition
Definition of Parmanu
🗒️ Major Postulates Of Kanad's Atomic Theory
  • All matter is composed of extremely small particles called Parmanu.
  • Parmanu are indivisible and cannot be further broken down.
  • Atoms are eternal and cannot be created or destroyed.
  • Different substances are formed by different combinations of atoms.
  • Atoms remain unchanged during physical and chemical transformations.
  • Changes in matter occur because atoms combine, separate, or rearrange themselves.
  • Atoms may exist in two states: motion and rest.
  • The properties of substances depend upon the arrangement and combination of atoms.
🗂️ Types / Category
Six Fundamental Categories (Padarthas)
Dravya
Substance or matter
Guna
Quality or property
Karma
Action or motion
Samanya
Universality or common characteristics
Visesa
Particularity or uniqueness
Samavaya
Inherence or inseparable relationship
🤔 Did You Know?
How Did Kanad Explain Formation of Matter?

According to Kanad, atoms combine in different ways to form larger particles and eventually visible matter.

For example:

  • Single atoms combine to form small clusters.
  • Several clusters combine to form larger particles.
  • Large numbers of particles produce observable matter.

Thus, different arrangements of the same basic particles can produce different substances.

⚖️ Comparison
Comparison with Modern Atomic Theory
Kanad's View Modern Scientific View
Matter consists of tiny particles called Parmanu. Matter consists of atoms.
Atoms are indivisible. Atoms contain electrons, protons, and neutrons.
Atoms combine to form substances. Atoms combine to form molecules and compounds.
Atoms are eternal. Atoms can undergo nuclear changes.
Based mainly on philosophical reasoning. Based on experimental evidence.
🏛️ Historical Note
Democritus and Leucippus

Around the same period, Greek philosophers Democritus and Leucippus proposed a similar idea. They suggested that matter could not be divided endlessly. Eventually, a smallest indivisible particle would be obtained.

Democritus called these particles Atomos, meaning indivisible. The modern term atom is derived from this Greek word.

⚖️ Kanad and Democritus: Similarities
  • Both believed matter is composed of tiny particles.
  • Both considered these particles indivisible.
  • Both proposed that different substances arise from different combinations of particles.
  • Neither theory was supported by experimental evidence.
🌟 Importance of Kanad's Atomic Theory
⚡ Exam Tip
❌ Common Mistakes
  • Confusing Parmanu with the modern atom.
  • Assuming Kanad's theory was experimentally proven.
  • Writing Democritus as the founder of modern atomic theory instead of Dalton.
  • Believing ancient atomic theories explained subatomic particles.
✏️ Example
Concept Check Question
Why is Kanad's atomic theory considered important even though it lacked experimental evidence?
  1. 1
    Identify the main idea proposed by Kanad.
  2. 2
    Connect it with modern atomic theory.
  3. 3
    Explain its historical significance.
Kanad was among the first thinkers to suggest that matter consists of extremely small indivisible particles called Parmanu. Although his ideas were philosophical and not experimentally verified, they introduced the fundamental concept of atomic structure. This concept later became the basis of modern atomic theories developed through scientific experiments.
📋 CBSE Competency-Based Case Study Question

A student reads that ancient Indian philosopher Kanad proposed that matter consists of tiny indivisible particles called Parmanu. He also learns that modern atoms contain electrons, protons, and neutrons.

Question 1: Which feature of Kanad's theory matches modern atomic theory?

Answer: Matter is composed of tiny particles.

Question 2: Which feature is not accepted today?

Answer: The idea that atoms are indivisible.

Question 3: Why is Kanad still remembered in science?

Answer: He introduced one of the earliest atomic concepts in history.

⚡ Quick Revision
  • Maharishi Kanad proposed the concept of Parmanu around 500 BCE.
  • Parmanu was considered the smallest indivisible particle.
  • Atoms combine to form different substances.
  • Kanad's theory was philosophical rather than experimental.
  • Democritus used the term Atomos meaning indivisible.
  • Modern atomic theory developed much later through scientific experiments.
🎨 SVG Diagram
Atomic Concept Proposed by Kanad
Atomic Concept of Rishi Kanad (c. 600 BCE) "All matter is composed of indivisible Anu (Atoms)" PARAMANU Single / Indivisible DVYANUKA Binary / Dyad (Combination of 2 Paramanu) TRASARENU Triad / Visible Matter (3 Dvyanuka = 6 Paramanu) "Paramanu combines in dyads (Dvyanuka) and triads (Trasarenu) to form perceptible matter."
⚛️

Lavoisier's Theory and the Law of Conservation of Mass

📋
Table of Contents
15 sections
🗺️ Overview
Antoine Lavoisier is widely regarded as the Father of Modern Chemistry. Through carefully designed quantitative experiments, he established one of the most important laws in chemistry: the Law of Conservation of Mass. This law states that matter can neither be created nor destroyed during a chemical reaction. Lavoisier's work transformed chemistry from a qualitative subject into a quantitative science based on precise measurements.
🖼️ Figure
Antoine Lavoisier Father of Modern Chemistry
Antoine Lavoisier (1743–1794), Father of Modern Chemistry
🏛️ Historical Note

Before Lavoisier's work, chemists often relied on observations rather than measurements. Many scientists believed in the Phlogiston Theory, which attempted to explain burning substances. Lavoisier conducted experiments using sealed containers and highly accurate balances. His observations proved that the total mass remains unchanged during chemical reactions.

Lavoisier introduced the practice of weighing reactants and products accurately, making chemistry a measurement-based science.
📘 Definition
Law of Conservation of Mass
🗒️ Lavoisier's Experimental Verification
Lavoisier's Experimental Verification

Lavoisier heated substances in sealed containers and measured their masses before and after the reaction.

  1. The mass of the sealed vessel and reactants was measured.
  2. The chemical reaction was allowed to occur.
  3. The vessel was weighed again.
  4. No change in total mass was observed.

This demonstrated that matter is conserved during chemical reactions.

🎨 SVG Diagram
Experimental Verification of Conservation of Mass
Conservation of Mass: Experimental Verification Before Reaction 150.00g After Reaction 150.00g Mass of Reactants = Mass of Products
🌟 Importance in Atomic Theory
✏️ Example
Solved Example
Hydrogen reacts with oxygen to form water. If 4 g of hydrogen reacts completely with 32 g of oxygen, what mass of water will be formed?
  1. 1
    Identify masses of reactants.
  2. 2
    Apply conservation of mass.
  3. 3
    Mass of products equals mass of reactants.
Given: Hydrogen = 4 g
Oxygen = 32 g
  1. Total mass of reactants:\[4 + 32 = 36\text{ g}\]
    According to the Law of Conservation of Mass:
    \[\text{Mass of water}=36\text{ g}\]
Answer: 36 g of water is formed.
✏️ Example
Examples from Everyday Life
  • Rusting of iron follows conservation of mass.
  • Burning of magnesium ribbon follows conservation of mass.
  • Formation of water from hydrogen and oxygen follows conservation of mass.
  • Photosynthesis and respiration also obey this law.
Apparent loss of mass is often observed because gases escape into the atmosphere. In a closed system, total mass remains constant.
🤔 Did You Know?
Why Does Mass Sometimes Appear to Change?

Students often observe that burning wood leaves only a small amount of ash. This may suggest a loss of mass, but actually gases such as carbon dioxide and water vapour escape into the air.

If all products are collected and measured, the total mass remains unchanged.

📘 Definition
Lavoisier's Definition of an Element
🗒️ Contribution
Contribution to Chemical Nomenclature

Lavoisier helped develop a systematic method of naming chemicals. Prior to this, chemical names were often confusing and inconsistent.

Modern chemical nomenclature originated from the system introduced by Lavoisier and his collaborators.

🗒️ Rejection
Rejection of the Phlogiston Theory

The Phlogiston Theory claimed that combustible materials released an invisible substance called phlogiston during burning.

Lavoisier demonstrated that combustion involves oxygen from air rather than the release of phlogiston. This discovery revolutionized chemistry and established the modern understanding of combustion.

⚡ Exam Tip
❌ Common Mistakes
Common Mistakes Made by Students
  • Writing that mass is conserved only in physical changes.
  • Confusing mass with volume.
  • Ignoring gaseous products while applying conservation of mass.
  • Assuming matter disappears during burning.
📋 CBSE Competency-Based Case Study

A student burns magnesium ribbon in a covered crucible. The mass of magnesium before heating is 2.4 g. After the reaction, 4.0 g of magnesium oxide is formed.

Question 1: Why is the mass of product greater than the mass of magnesium?

Answer: Oxygen from the air combines with magnesium.

Question 2: Does this violate conservation of mass?

Answer: No. The mass of oxygen added must also be included.

Question 3: Which law explains this observation?

Answer: Law of Conservation of Mass.

⚡ Quick Revision
  • Antoine Lavoisier is called the Father of Modern Chemistry.
  • He proposed the Law of Conservation of Mass.
  • Total mass of reactants equals total mass of products.
  • Atoms are neither created nor destroyed during chemical reactions.
  • He rejected the Phlogiston Theory.
  • He helped define elements and modern chemical nomenclature.
⚛️

Laws of Chemical Combination

📋
Table of Contents
14 sections
📘 Definition
🌟 Why are the Laws of Chemical Combination Important?
⚖️ Laws
Major Laws of Chemical Combination
Law Scientist Main Idea
Law of Conservation of Mass Antoine Lavoisier Mass is neither created nor destroyed during a chemical reaction.
Law of Definite Proportions Joseph Proust A compound always contains the same elements in a fixed ratio by mass.
⚖️ Law of Conservation of Mass

Mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of reactants is always equal to the total mass of products.

Mathematical Representation

\[ \text{Total Mass of Reactants} = \text{Total Mass of Products} \]

\[ m_r = m_p \]

Explanation

During a chemical reaction, atoms are neither created nor destroyed. They only rearrange themselves to form new substances. Since the number of atoms remains unchanged, total mass also remains constant.

Experimental Verification

Lavoisier performed chemical reactions in sealed containers and measured masses before and after the reaction. He found that the total mass remained unchanged, leading to the formulation of this law.

REACTANTS Total Mass: 50 g Chemical Reaction PRODUCTS Total Mass: 50 g Mass Before Reaction = Mass After Reaction
✏️ Example
Solved Example
10 g of calcium reacts with 18 g of water. Calculate the total mass of products formed.
  1. 1
    Find total mass of reactants.
  2. 2
    Apply conservation of mass.
  3. 3
    Mass of products equals mass of reactants.
  1. Mass of Reactants = 10 + 18 = 28 g
    According to the Law of Conservation of Mass:
    \[\text{Mass of Products} = 28\text{ g}\]
28 g
⚖️ Law of Definite Proportions (Law of Constant Proportions)
Scientist

This law was proposed by Joseph Proust in 1799.

Statement

A pure chemical compound always contains the same elements combined together in the same fixed proportion by mass, irrespective of its source or method of preparation.

No matter where a compound is found or how it is prepared, the ratio of masses of its constituent elements remains constant.

For example, water obtained from:

  • Rain
  • River
  • Sea
  • Laboratory preparation

always contains hydrogen and oxygen in the same mass ratio.

bg-dark bg-gradient p-2 rounded

Water consists of hydrogen and oxygen.

\[ \text{Mass of Hydrogen} : \text{Mass of Oxygen} = 1:8 \]

Thus, every 1 g of hydrogen combines with 8 g of oxygen to form water.

Example of Carbon Dioxide

Carbon dioxide always contains carbon and oxygen in a fixed mass ratio.

\[ \text{Carbon} : \text{Oxygen} = 3:8 \]

This ratio remains unchanged irrespective of the source of carbon dioxide.

Illustration of Constant Composition
Sample A River Water H : O = 1 : 8 IDENTICAL Constant Composition Sample B Ocean Water H : O = 1 : 8 Law of Constant Proportions Elements are always present in definite proportions by mass.
✏️ Example

Water obtained from different sources contains:

  • 11.1% hydrogen by mass
  • 88.9% oxygen by mass

Therefore, all pure samples of water have identical composition.

✏️ Example
Solved Example
A sample of water contains 2 g hydrogen and 16 g oxygen. Does it obey the Law of Definite Proportions?
\[2:16 = 1:8\]

Since the mass ratio remains 1:8, the sample obeys the Law of Definite Proportions.

⚖️ Comparison
Difference Between the Two Laws
Law of Conservation of Mass Law of Definite Proportions
Concerned with total mass during a reaction. Concerned with composition of a compound.
Proposed by Lavoisier. Proposed by Proust.
Mass remains constant. Mass ratio remains constant.
Applicable to chemical reactions. Applicable to chemical compounds.
🗒️ Connection with Dalton's Atomic Theory

Dalton proposed his Atomic Theory to explain these laws.

  • Atoms cannot be created or destroyed → explains Conservation of Mass.
  • Atoms combine in fixed whole-number ratios → explains Definite Proportions.
  • Compounds are formed by combination of atoms in simple ratios.
⚡ Exam Tip
❌ Common Mistakes
  • Confusing Conservation of Mass with Constant Proportions.
  • Writing volume ratio instead of mass ratio.
  • Ignoring gaseous reactants or products in mass calculations.
  • Assuming compounds may have variable composition.
📋 CBSE Competency-Based Case Study

Two students prepare water in different laboratories. One obtains water from river water purification, while the other prepares it by combining hydrogen and oxygen in a laboratory.

Analysis shows that both samples contain hydrogen and oxygen in the mass ratio 1:8.

Question 1: Which law is illustrated?

Answer: Law of Definite Proportions.

Question 2: Why is the composition identical?

Answer: Pure compounds always contain elements in fixed mass ratios.

Question 3: Which scientist proposed this law?

Answer: Joseph Proust.

⚡ Quick Revision
  • Laws of Chemical Combination form the basis of modern chemistry.
  • Lavoisier proposed the Law of Conservation of Mass.
  • Proust proposed the Law of Definite Proportions.
  • Total mass remains constant during chemical reactions.
  • Compounds always have fixed composition by mass.
  • These laws provided evidence for Dalton's Atomic Theory.
⚛️

Dalton's Atomic Theory

📋
Table of Contents
13 sections
🖼️ Figure
John Dalton (1766–1844), the scientist who proposed the first modern atomic theory
John Dalton Atomic Theory
🗺️ Overview

In 1808, John Dalton proposed the first scientifically accepted atomic theory to explain the laws of chemical combination. Dalton's theory transformed chemistry by suggesting that all matter is composed of tiny particles called atoms. His theory provided a logical explanation for the Law of Conservation of Mass and the Law of Definite Proportions.

🏛️ Historical Note

Before Dalton's work, scientists had already discovered several important laws of chemical combination. However, no satisfactory explanation existed for why elements combined in fixed ratios or why mass remained conserved during chemical reactions.

Dalton proposed that matter consists of tiny indivisible particles called atoms, and this idea successfully explained the experimental observations of his time.

Dalton's Atomic Theory was the first theory to explain chemical reactions using particles of matter.
🤔 Did You Know?
Why Was Dalton's Atomic Theory Needed?
  • To explain the Law of Conservation of Mass.
  • To explain the Law of Definite Proportions.
  • To understand how compounds are formed.
  • To provide a scientific model for matter.
  • To explain chemical reactions at the particle level.
📘 Definition
Definition of Dalton's Atomic Theory
🗂️ Postulates of Dalton's Atomic Theory
Postulate 1: Matter is Made of Atoms
All matter, whether an element, compound, or mixture, is made up of extremely tiny particles called atoms.
These atoms are the basic building blocks of matter.
Postulate 2: Atoms are Indivisible
Dalton believed that atoms cannot be divided into smaller particles and cannot be created or destroyed during chemical reactions.
This postulate explained the Law of Conservation of Mass.
Postulate 3: Atoms of the Same Element are Identical
All atoms of a particular element have identical masses and similar chemical properties.
Example:
  • All oxygen atoms were believed to be identical.
  • All hydrogen atoms were believed to be identical.
Postulate 4: Atoms of Different Elements Differ
Atoms of different elements possess different masses and chemical properties.
Example:
  • Hydrogen atoms differ from oxygen atoms.
  • Carbon atoms differ from nitrogen atoms.
Postulate 5: Atoms Combine in Simple Whole-Number Ratios
Atoms combine in simple whole-number ratios to form compounds.
Examples:
  • Water (\(H_2O\)) contains 2 hydrogen atoms and 1 oxygen atom.
  • Carbon dioxide (\(CO_2\)) contains 1 carbon atom and 2 oxygen atoms.
  • Ammonia (\(NH_3\)) contains 1 nitrogen atom and 3 hydrogen atoms.
Postulate 6: Composition of a Compound Remains Fixed
The relative number and kinds of atoms remain constant in a given compound.
Therefore, water always contains hydrogen and oxygen in the same proportion.
🎨 SVG Diagram
Visual Representation of Dalton's Atomic Theory
Dalton's Atomic Theory 1. INDIVISIBLE ATOMS Solid, Hard Spheres 2. IDENTICAL MASS Same Mass & Properties 3. COMPOUND FORMATION Fixed Ratios by Mass 4. CHEMICAL REACTION AS REARRANGEMENT REACTANTS PRODUCTS Atoms are neither created nor destroyed, only rearranged.
🔍 How Dalton's Theory Explained the Laws of Chemical Combination
Explanation of Conservation of Mass

Dalton proposed that atoms cannot be created or destroyed during chemical reactions. Therefore, total mass remains constant.

Explanation of Definite Proportions

Dalton stated that compounds are formed by fixed combinations of atoms. Hence, compounds always have a constant composition.

Examples Based on Dalton's Atomic Theory
Example 1: Formation of Water

Water contains hydrogen and oxygen atoms in a fixed ratio:

\[ H_2O \]

Every water molecule contains two hydrogen atoms and one oxygen atom.

Example 2: Formation of Carbon Dioxide

\[ CO_2 \]

Every carbon dioxide molecule contains one carbon atom and two oxygen atoms.

⚠️ Limitations of Dalton's Atomic Theory

Although Dalton's theory was revolutionary, later discoveries showed that some of its postulates were not entirely correct.

Dalton's Postulate Modern Discovery
Atoms are indivisible. Atoms contain electrons, protons, and neutrons.
Atoms cannot be created or destroyed. Nuclear reactions can create or destroy atoms.
All atoms of an element are identical. Isotopes of the same element have different masses.
Atoms of different elements always differ in mass. Different elements may have atoms with the same mass number.
📌 Modern View of the Atom
⚡ Exam Tip
❌ Common Mistakes
  • Writing Dalton instead of Lavoisier for Conservation of Mass.
  • Confusing atoms with molecules.
  • Memorising postulates without understanding their significance.
  • Ignoring the limitations of Dalton's theory.
  • Writing modern atomic concepts as Dalton's original ideas.
📋 CBSE Competency-Based Case Study

A student observes that water always contains hydrogen and oxygen in a fixed composition. He also notices that the total mass remains unchanged when hydrogen reacts with oxygen to form water.

Question 1: Which theory explains these observations?

Answer: Dalton's Atomic Theory.

Question 2: Which law is related to fixed composition?

Answer: Law of Definite Proportions.

Question 3: Which postulate explains conservation of mass?

Answer: Atoms cannot be created or destroyed during chemical reactions.

🗒️ Assertion and Reason Practice

Assertion (A): Water always contains hydrogen and oxygen in a fixed ratio.

Reason (R): Atoms combine in simple whole-number ratios.

Answer: Both Assertion and Reason are true, and Reason correctly explains Assertion.

⚡ Quick Revision
  • Dalton proposed the first modern atomic theory in 1808.
  • All matter is made of atoms.
  • Atoms combine in simple whole-number ratios.
  • Compounds possess fixed composition.
  • Dalton explained the laws of chemical combination.
  • Modern discoveries revealed certain limitations in his theory.
⚛️

Atom – The Fundamental Unit of Matter

📋
Table of Contents
14 sections
🗺️ Overview
An atom is the smallest particle of an element that retains the chemical properties of that element. Atoms are the basic building blocks of all matter. Every substance around us, whether solid, liquid, or gas, is ultimately made up of atoms.
The word Atom is derived from the Greek word "Atomos", meaning indivisible.
📘 Definition
🌟 Why are Atoms Important?
🔷 Characteristics of Atoms
🔷 Characteristics
  • Atoms are extremely small particles.
  • Atoms cannot be seen with the naked eye.
  • Atoms of different elements differ in mass and properties.
  • Atoms participate in chemical reactions.
  • Atoms generally combine to form molecules.
  • Atoms retain the chemical identity of their element.
🤔 Did You Know?
How Small is an Atom?
Atoms are extraordinarily small. Their size is measured in nanometres (nm).
\[1\ \text{nm}=10^{-9}\ \text{m}\]
Typical atomic sizes are of the order:
\[10^{-10}\ \text{m}\]
To appreciate this size:
  • A sheet of paper contains billions of atoms.
  • A human hair is approximately one million carbon atoms thick.
  • Atoms are far too small to be observed directly with ordinary microscopes.
⚖️ Comparison of Sizes
Object Approximate Size
Football \(10^{-1}\) m
Ant \(10^{-2}\) m
Human Hair Thickness \(10^{-4}\) m
Bacterium \(10^{-6}\) m
Atom \(10^{-10}\) m
🎨 SVG Diagram
Visualising the Tiny Size of Atoms
Visualizing the Tiny Size of Atoms 1 mm Visible Tip of a Pin 1,000 nm Typical Bacterium 0.1 nm Radius of an Atom An atom is 10,000,000 times smaller than a millimeter.
📌 Symbols of Elements
📌 Note
Rules for Writing Symbols
✏️ Example
  • Hydrogen → H
  • Carbon → C
  • Helium → He
  • Chlorine → Cl
Symbols Derived from Latin Names
Element Symbol Latin Name
Sodium Na Natrium
Potassium K Kalium
Iron Fe Ferrum
Copper Cu Cuprum
Silver Ag Argentum
Gold Au Aurum
Mercury Hg Hydrargyrum
🤔 Did You Know?
Why Do Most Atoms Not Exist Independently?
Most atoms are highly reactive and therefore combine with other atoms to achieve stability.
As a result, atoms usually exist in the form of:
  • Molecules
  • Compounds
  • Ions
Only a few elements such as helium, neon, argon, krypton and xenon exist as single atoms under normal conditions.
⚡ Exam Tip
❌ Common Mistakes
  • Writing the second letter of a symbol in capital form (e.g., CL instead of Cl).
  • Confusing symbols derived from Latin names.
  • Assuming all atoms can exist independently.
  • Confusing atoms with molecules.
  • Using lowercase letters for the first letter of a symbol.
📋 CBSE Competency-Based Case Study

A student writes the symbols of sodium, potassium and iron as S, P and I respectively.

Question 1: Are these symbols correct?

Answer: No.

Question 2: Write the correct symbols.

Answer: Na, K and Fe.

Question 3: Why are these symbols different from English names?

Answer: They are derived from Latin names.

⚡ Quick Revision
  • Atom is the basic unit of an element.
  • Atoms are extremely small particles.
  • Symbols are standardized by IUPAC.
  • The first letter of a symbol is always capital.
  • Many symbols are derived from Latin names.
  • Most atoms do not exist independently.
⚛️

IUPAC (International Union of Pure and Applied Chemistry)

📋
Table of Contents
3 sections
📖 Introduction
🌟 Major Functions of IUPAC
🤔 Did You Know?
Why is IUPAC Important?
  • Ensures uniformity in scientific communication.
  • Prevents confusion caused by local names of chemicals.
  • Allows scientists from different countries to understand one another.
  • Provides globally accepted standards for chemistry education.
  • Maintains consistency in textbooks, research papers and laboratories.
⚛️

Atomic Mass

📋
Table of Contents
10 sections
🗺️ Overview
Atoms are extremely small particles, and their actual masses are extraordinarily tiny. Measuring atomic masses directly in grams is impractical. Therefore, scientists use a special unit called the atomic mass unit (u) to express atomic masses conveniently.
🌟 Why Do We Need Atomic Mass?
📘 Definition
🤔 Did You Know?
Why Was Carbon-12 Chosen as the Standard?
  • It is stable.
  • It is readily available.
  • It gives convenient atomic mass values.
  • It is accepted internationally.
📌 Understanding Relative Atomic Mass
🎨 SVG Diagram
Visual Comparison of Atomic Masses
RELATIVE ATOMIC MASS A Comparison of Mass Units (u) based on Carbon-12 STEP 1: THE STANDARD Carbon-12 Atom 1 unit (u) = 1/12th Mass of Carbon-12 STEP 2: THE COMPARISON Oxygen (O) 16 Units (u) MASS = 16 u Relative to 1/12th of Carbon-12 Atomic mass is not absolute weight, but a ratio relative to the standard Carbon-12 unit.
ℹ️ Atomic Masses of Important Elements
Element Symbol Atomic Mass (u)
Hydrogen H 1
Carbon C 12
Nitrogen N 14
Oxygen O 16
Sodium Na 23
Magnesium Mg 24
Sulphur S 32
Chlorine Cl 35.5
Calcium Ca 40
✏️ Example
Solved Example
Which atom is heavier: Oxygen or Carbon?
  1. Atomic mass of Carbon:\[12u\]
    Atomic mass of Oxygen:
    \[16u\]
  2. Since \[16 \gt 12\]
    therefore, Oxygen is heavier than Carbon.
❌ Common Mistakes
  • Confusing atomic mass with atomic number.
  • Writing atomic mass unit as "gm".
  • Assuming atomic mass means exact mass in grams.
  • Forgetting that Carbon-12 is the standard reference.
  • Confusing atomic mass with molecular mass.
📋 CBSE Competency-Based Question

A student states that oxygen has an atomic mass of 16 u while hydrogen has an atomic mass of 1 u.

Question: What does this indicate?

Answer: Oxygen atoms are approximately sixteen times heavier than one atomic mass unit and much heavier than hydrogen atoms.

⚡ Quick Revision
  • Atomic masses are expressed in atomic mass units (u).
  • Carbon-12 is the international standard.
  • \(1u\) equals one-twelfth of the mass of a Carbon-12 atom.
  • Atomic mass represents relative mass.
  • Atomic mass should not be confused with atomic number.
⚛️

Molecules

📋
Table of Contents
13 sections
🗺️ Overview
A molecule is a group of two or more atoms chemically bonded together and held by attractive forces. It is the smallest particle of an element or compound that can exist independently and retains all the chemical properties of that substance.
A molecule behaves as a single unit and possesses all the characteristics of the substance to which it belongs.
📘 Definition
🤔 Did You Know?
Why Do Molecules Form?
Most atoms cannot exist independently because they are unstable. To attain stability, atoms combine with other atoms and form molecules. During this process, atoms become more stable and attain lower energy.
Examples
  • Two hydrogen atoms combine to form a hydrogen molecule (\(H_2\)).
  • Two oxygen atoms combine to form an oxygen molecule (\(O_2\)).
  • One carbon atom combines with two oxygen atoms to form carbon dioxide (\(CO_2\)).
🔷 Characteristics of Molecules
🔷 Characteristics
  • Molecules consist of two or more atoms chemically bonded together.
  • Molecules are electrically neutral.
  • They can exist independently under suitable conditions.
  • They retain all properties of the substance.
  • Atoms combine in definite ratios to form molecules.
🗂️ Types of Molecules
Molecules of Elements
Molecules consisting of atoms of the same element are called molecules of elements.
Examples
Element Molecular Formula Number of Atoms
Hydrogen \(H_2\) 2
Oxygen \(O_2\) 2
Nitrogen \(N_2\) 2
Ozone \(O_3\) 3
Phosphorus \(P_4\) 4
Sulphur \(S_8\) 8
Molecules of Compounds
Molecules consisting of atoms of different elements combined in fixed proportions are called molecules of compounds.
Examples
Compound Formula Constituent Atoms
Water \(H_2O\) Hydrogen and Oxygen
Carbon Dioxide \(CO_2\) Carbon and Oxygen
Ammonia \(NH_3\) Nitrogen and Hydrogen
Methane \(CH_4\) Carbon and Hydrogen
📘 Definition
Atomicity
🗒️ Classification Of Molecules Based On Atomicity
Type Number of Atoms Examples
Monoatomic 1 He, Ne, Ar
Diatomic 2 H₂, O₂, N₂, Cl₂
Triatomic 3 O₃, CO₂
Tetra-atomic 4 P₄, NH₃ (4 atoms total)
Polyatomic More than 4 S₈, C₆H₁₂O₆
⚖️ Difference Between Atom and Molecule
Atom Molecule
Smallest particle of an element. Smallest particle capable of independent existence.
May or may not exist independently. Can exist independently.
Single particle. Contains two or more atoms.
Example: H, O, Na Example: H₂, O₂, H₂O
✏️ Example
Solved Example
Determine the atomicity of ozone (\(O_3\)).
  1. 1
    Identify the number of atoms present.
  2. 2
    Count the subscript.
  3. 3
    State atomicity.

Ozone is represented by:

\[ O_3 \]

Therefore, one molecule contains 3 oxygen atoms.

Atomicity = 3 (Triatomic)

⚡ Exam Tip
❌ Common Mistakes
  • Confusing atomicity with molecular mass.
  • Writing O₃ as a diatomic molecule.
  • Assuming all molecules contain atoms of different elements.
  • Confusing molecules with compounds.
  • Ignoring the concept of independent existence.
📋 CBSE Competency-Based Question

A student observes the molecules \(H_2\), \(O_3\), \(CO_2\), and \(P_4\).

Question 1: Which molecule is triatomic?

Answer: \(O_3\).

Question 2: Which molecule contains atoms of different elements?

Answer: \(CO_2\).

Question 3: What is the atomicity of \(P_4\)?

Answer: 4 (Tetra-atomic).

⚡ Quick Revision
  • Molecules are formed by chemical combination of atoms.
  • Molecules can be of elements or compounds.
  • Atomicity is the number of atoms present in one molecule.
  • H₂ and O₂ are diatomic molecules.
  • O₃ is triatomic.
  • S₈ is polyatomic.
⚛️

Ions

📋
Table of Contents
13 sections
🗺️ Overview
An ion is an atom or a group of atoms that carries a net electric charge due to the loss or gain of one or more electrons. Ions are formed when the number of electrons becomes unequal to the number of protons present in an atom or group of atoms.
Atoms are electrically neutral, whereas ions carry positive or negative charges.
📘 Definition
Definition of an Ion
🤔 Did You Know?
Why Do Ions Form?
Atoms tend to attain a more stable electronic configuration. To achieve stability, atoms may lose or gain electrons and form ions.
  • Loss of electrons → Positive ion (cation)
  • Gain of electrons → Negative ion (anion)
🗂️ Formation of Ions
Formation of a Positive Ion
When sodium loses one electron:
\[Na \rightarrow Na^{+} + e^{-}\]
Sodium becomes a positively charged ion because it has lost one negative electron.
Formation of a Negative Ion
When chlorine gains one electron:
\[Cl + e^{-} \rightarrow Cl^{-}\]
Chlorine becomes a negatively charged ion because it gains an extra electron.
🗂️ Types of Ions
Cations
Positively charged ions formed by the loss of electrons are called cations.
Loss of electrons → Positive charge → Cation
Examples
Atom Ion Formed Charge
Sodium \(Na^+\) +1
Potassium \(K^+\) +1
Calcium \(Ca^{2+}\) +2
Magnesium \(Mg^{2+}\) +2
Aluminium \(Al^{3+}\) +3
Anions
Negatively charged ions formed by the gain of electrons are called anions.
Gain of electrons → Negative charge → Anion
Examples
Atom Ion Formed Charge
Chlorine \(Cl^-\) -1
Fluorine \(F^-\) -1
Oxygen \(O^{2-}\) -2
Sulphur \(S^{2-}\) -2
Nitrogen \(N^{3-}\) -3
📌 Classification of Ions
Ions may also be classified according to the number of atoms present.
Monoatomic Ions
Ions consisting of a single atom carrying a charge are called monoatomic ions.
Examples
Ion Name Formula Charge
Sodium \(Na^+\) +1
Potassium \(K^+\) +1
Magnesium \(Mg^{2+}\) +2
Chloride \(Cl^-\) -1
Oxide \(O^{2-}\) -2
Polyatomic Ions
Ions consisting of two or more atoms bonded together and carrying a net charge are called polyatomic ions.
Examples
Ion Name Formula Charge
Ammonium \(NH_4^+\) +1
Hydroxide \(OH^-\) -1
Nitrate \(NO_3^-\) -1
Carbonate \(CO_3^{2-}\) -2
Sulphate \(SO_4^{2-}\) -2
⚖️ Difference Between Atom and Ion
Atom Ion
Electrically neutral. Electrically charged.
Equal number of protons and electrons. Unequal number of protons and electrons.
No net charge. Positive or negative charge.
Example: Na, Cl Example: Na⁺, Cl⁻
🌟 Importance of Ions
✏️ Example
Solved Example
Identify whether Mg2+ is a cation or an anion.
Mg2+ carries a positive charge.
Positively charged ions are called cations.
Therefore, Mg2+ is a cation.
⚡ Exam Tip
❌ Common Mistakes
  • Confusing cations with anions.
  • Writing chloride ion as \(Cl^+\).
  • Forgetting charges while writing polyatomic ions.
  • Assuming all ions consist of single atoms.
  • Confusing ions with molecules.
📋 CBSE Competency-Based Question

A student is given the species: \(Na^+\), \(Cl^-\), \(NH_4^+\), \(SO_4^{2-}\).

Classify them as cations and anions.

Solution

Cations: \(Na^+\), \(NH_4^+\)

Anions: \(Cl^-\), \(SO_4^{2-}\)

⚡ Quick Revision
  • Ions are charged particles formed by loss or gain of electrons.
  • Cations are positively charged ions.
  • Anions are negatively charged ions.
  • Monoatomic ions contain one atom.
  • Polyatomic ions contain more than one atom.
  • Ions are essential for writing chemical formulae and understanding valency.
⚛️

Valency

📋
Table of Contents
12 sections
🗺️ Overview
Valency is the combining capacity of an atom or a group of atoms. It indicates the number of electrons an atom loses, gains, or shares while forming a chemical bond. Valency helps us understand how atoms combine to form molecules and compounds.

Valency is one of the most important concepts in chemistry because it forms the basis for writing chemical formulae and understanding the formation of compounds.
📘 Definition
🌟 Why is Valency Important?
🤔 Did You Know?
How Does Valency Arise?
Atoms combine with one another to attain a stable electronic arrangement. During this process, atoms may lose, gain, or share electrons.
The number of electrons involved in this process determines the valency of the atom.
For Class 9, remember:
Valency is mainly the combining capacity of an atom and is often equal to the charge on the ion formed by that atom.
📌 Determination of Valency
🎨 SVG Diagram
Visual Understanding of Valency
Visualizing Valency: Combining Capacity Valency = 1 Hydrogen (H) Valency = 2 Oxygen (O) Valency = 4 Carbon (C) The "Lock & Key" Bonding Concept Methane (CH4) Carbon uses all 4 "hooks" to bond with 4 Hydrogen atoms. Every hook must be filled. Valency is the number of electrons an atom needs to gain, lose, or share to reach a stable state.
💡 Concept
Using Valency to Write Chemical Formulae
⚖️ Difference Between Valency and Charge
Valency Charge
Combining capacity of an atom. Electrical property of an ion.
Always expressed as a number. Expressed as positive or negative sign.
Example: Oxygen = 2 Example: \(O^{2-}\)
✏️ Example
Solved Examples
What is the valency of oxygen?
Oxygen generally combines with two hydrogen atoms in water.
Therefore:
\[\text{Valency of Oxygen} = 2\]
What is the valency of aluminium?
Aluminium forms the ion: Al3+
Therefore:
Valency of Aluminium = 3
⚡ Exam Tip
❌ Common Mistakes
  • Confusing valency with atomicity.
  • Confusing valency with charge.
  • Writing valency of oxygen as 1.
  • Ignoring radicals while writing formulae.
  • Using charge signs instead of valencies during formula writing.
📋 CBSE Competency-Based Question

A student is given the following valencies:

  • Aluminium = 3
  • Oxygen = 2

Write the formula of the compound formed.

Solution

Applying the criss-cross method:

\[ Al_2O_3 \]

Therefore, the compound formed is aluminium oxide.

⚡ Quick Revision
  • Valency is the combining capacity of an atom or radical.
  • Valency helps in writing chemical formulae.
  • Oxygen has valency 2.
  • Carbon has valency 4.
  • Sodium has valency 1.
  • Valency should not be confused with atomicity or charge.
⚛️

Chemical Formula

📋
Table of Contents
14 sections
🗺️ Overview
A chemical formula is the symbolic representation of a substance that shows the elements present and the number of atoms of each element in one molecule or formula unit of the substance.
A chemical formula provides concise information about the composition of a compound without writing the full names of its constituent elements.
📘 Definition
efinition of Chemical Formula
🌟 Importance of Chemical Formulae
⚖️ Rules for Writing Chemical Formulae
  1. The total positive charge must balance the total negative charge.
  2. In compounds formed from a metal and a non-metal, the symbol of the metal is written first.
  3. In compounds containing a positive ion and a negative ion, the positive ion is written first.
  4. The valencies of the combining species are used to determine the formula.
  5. For polyatomic ions, brackets are used when more than one ion is present.
  6. If only one polyatomic ion is present, brackets are not required.
Examples
  • \(NaCl\)
  • \(CaO\)
  • \(Mg(OH)_2\)
  • \(Al_2(SO_4)_3\)
🔄 Steps for Writing Chemical Formulae
  • 1
    Write the symbols of the constituent ions or elements.
  • 2
    Write their valencies.
  • 3
    Criss-cross the valencies.
  • 4
    Write the crossed numbers as subscripts.
  • 5
    Simplify the ratio if possible.
  • 6
    Check that total positive and negative charges are balanced.
📘 The Criss-Cross Method
✏️ Writing Formulae Involving Polyatomic Ions
When a compound contains more than one polyatomic ion, brackets must be used.
Magnesium ion: Mg2+
Hydroxide ion: OH-
Applying the criss-cross method: Mg(OH)2
🌟 Important Valencies for Formula Writing
✏️ Example
Solved Example
Write the formula of calcium hydroxide.
Calcium = 2
Hydroxide = 1
Ca(OH)2
Write the formula of sodium oxide.

Sodium = 1
Oxygen = 2

Na2O

⚖️ Difference Between Symbol and Chemical Formula
Symbol Chemical Formula
Represents one element. Represents a compound or molecule.
Example: H, O, Na Example: H₂O, CO₂, NaCl
Represents one atom. Represents a molecule or formula unit.
⚡ Exam Tip
❌ Common Mistakes
  • Writing metal symbols after non-metals.
  • Forgetting to simplify valency ratios.
  • Not using brackets with polyatomic ions.
  • Writing \(MgOH_2\) instead of \(Mg(OH)_2\).
  • Confusing valency with atomicity.
📋 CBSE Competency-Based Question

A student writes the formula of aluminium sulphate as:

\[ AlSO_4 \]

Is the formula correct?

Solution

Aluminium has valency 3 and sulphate has valency 2.

Applying the criss-cross method:

\[ Al_2(SO_4)_3 \]

Therefore, \(AlSO_4\) is incorrect.

⚡ Quick Revision
  • A chemical formula represents the composition of a substance.
  • The total positive and negative charges must balance.
  • Valencies are used to determine subscripts.
  • The criss-cross method helps write formulae easily.
  • Brackets are required for multiple polyatomic ions.
  • Formula writing is one of the most important CBSE examination skills.
⚛️

Formulae of Simple Compounds

📋
Table of Contents
10 sections
📘 Definition
🤔 Did You Know?
What are Binary Compounds?
Binary compounds are compounds formed by the combination of only two different elements.
Examples
  • \(HCl\) — Hydrogen chloride
  • \(H_2S\) — Hydrogen sulphide
  • \(NaCl\) — Sodium chloride
  • \(MgO\) — Magnesium oxide
  • \(Al_2O_3\) — Aluminium oxide
🔄 General Steps for Writing Chemical Formulae
  • 1
    Write the symbols of the combining elements or ions.
  • 2
    Write their valencies.
  • 3
    Criss-cross the valencies.
  • 4
    Use the crossed numbers as subscripts.
  • 5
    Reduce the ratio if possible.
  • 6
    Verify that the total positive and negative valencies balance.
✏️ Example
Formula of Hydrogen Chloride
Hydrogen and chlorine combine in a 1:1 ratio because both have valency 1.
Element Symbol Valency
Hydrogen H 1
Chlorine Cl 1
Formula of Hydrogen Chloride
Hydrogen chloride contains one hydrogen atom and one chlorine atom.
\[HCl\]
✏️ Example
Formula of Hydrogen Sulphide
Hydrogen has valency 1 and sulphur has valency 2.
Element Valency
Hydrogen (H) 1
Sulphur (S) 2
Formula of Hydrogen Sulphide
Hydrogen sulphide contains two hydrogen atoms and one sulphur atom.
\[H_2S\]
✏️ Example
Formula of Aluminium Oxide
Aluminium has valency 3 while oxygen has valency 2.
Element Valency
Aluminium (Al) 3
Oxygen (O) 2
Formula of Aluminium Oxide
Aluminium oxide contains two aluminium atoms and three oxygen atoms.

Criss-Cross Method

\[ Al_2O_3 \]

✏️ Example
Formula of Ammonium Sulphate
Ammonium (\(NH_4^+\)) and sulphate (\(SO_4^{2-}\)) are polyatomic ions.
Ion Valency
\(NH_4^+\) 1
\(SO_4^{2-}\) 2
Formula of Ammonium Sulphate
Brackets are necessary because two ammonium ions are present.

Criss-Cross Method

\[ (NH_4)_2SO_4 \]

❌ Common Mistakes
  • Writing the non-metal before the metal.
  • Forgetting to simplify valencies.
  • Ignoring brackets for polyatomic ions.
  • Using charges instead of valencies directly.
  • Writing incorrect subscripts after criss-crossing.
📋 CBSE Competency-Based Question

A student writes the formula of calcium hydroxide as:

\[ CaOH_2 \]

Is the formula correct? Justify your answer.

Solution

Hydroxide is a polyatomic ion (\(OH^-\)).

Since two hydroxide ions are required, brackets must be used.

\[ Ca(OH)_2 \]

Therefore, the student's answer is incorrect.

⚡ Quick Revision
  • Binary compounds contain only two elements.
  • Chemical formulae are written using valencies.
  • The criss-cross method simplifies formula writing.
  • Positive ion is written before negative ion.
  • Brackets are used for multiple polyatomic ions.
  • Always verify that total positive and negative valencies balance.
⚛️

Molecular Mass

📋
Table of Contents
12 sections
🗺️ Overview

The molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance. It represents the relative mass of a molecule and is expressed in atomic mass units (u).

Molecular mass tells us how heavy a molecule is compared to one atomic mass unit (u).
📘 Definition
🔢 Formula

Formula for Molecular Mass

Molecular mass is calculated using:
\[\text{Molecular Mass} = \sum (\text{Atomic Mass} \times \text{Number of Atoms})\]
In simple words:
\[\text{Molecular Mass} = \text{Total of all atomic masses present in a molecule}\]
🌟 Why is Molecular Mass Important?
🔄 Steps to Calculate Molecular Mass
  • 1
    Write the chemical formula.
  • 2
    Identify the number of atoms of each element.
  • 3
    Write their atomic masses.
  • 4
    Multiply atomic mass by the number of atoms.
  • 5
    Add all the values.
🎨 SVG Diagram
Visual Understanding of Molecular Mass
Understanding Molecular Mass Example: Carbon Dioxide (CO₂) C Carbon: 12.01 u O Oxygen: 16.00 u O Oxygen: 16.00 u Molecular Mass Calculation 12.01 + (2 × 16.00) = 44.01 u Molecular Mass = Σ Atomic Masses
✏️ Example
Find Molecular Mass of Water (H₂O)
  1. 1
    Water contains:
    • 2 Hydrogen atoms
    • 1 Oxygen atom
Given: Atomic mass of Hydrogen = 1 u
Atomic mass of Oxygen = 16 u
  1. \[ \begin{aligned}\text{Molecular Mass of H}_2O &=(2 \times 1)+(1 \times 16)\\ &= 2 + 16\\ &=18 \end{aligned} \]
Molecular mass of water = 18 u.
Calculate Molecular Mass of Nitric Acid (HNO₃)
  1. 1
    One molecule of nitric acid contains:
    • 1 Hydrogen atom
    • 1 Nitrogen atom
    • 3 Oxygen atoms
Given:

H = 1 u
N = 14 u
O = 16 u

  1. \( (1\times1) + (1\times14) + (3\times16) \)
  2. =1+14+48
  3. =63u
Molecular mass of nitric acid = 63 u.
ℹ️ Molecular Masses of Common Compounds
Compound Formula Molecular Mass (u)
Water \(H_2O\) 18
Carbon Dioxide \(CO_2\) 44
Ammonia \(NH_3\) 17
Methane \(CH_4\) 16
Nitric Acid \(HNO_3\) 63
Glucose \(C_6H_{12}O_6\) 180
⚖️ Comparison
Atomic Mass Molecular Mass
Mass of one atom. Mass of one molecule.
Applicable to elements. Applicable to molecules.
Example: O = 16 u. Example: O₂ = 32 u.
⚡ Exam Tip
❌ Common Mistakes
  • Ignoring subscripts while calculating mass.
  • Using atomic number instead of atomic mass.
  • Adding atomic masses incorrectly.
  • Confusing molecular mass with formula unit mass.
  • Not counting atoms inside brackets.
📋 CBSE Competency-Based Question

A student calculates the molecular mass of carbon dioxide as:

\[ 12 + 16 = 28u \]

Identify the mistake.

Solution

Carbon dioxide contains two oxygen atoms.

\[ 12 + (2\times16) \]

\[ =44u \]

The student ignored the subscript 2.

⚡ Quick Revision
  • Molecular mass = Sum of atomic masses of all atoms in a molecule.
  • Unit of molecular mass is u.
  • Always multiply atomic mass by the number of atoms.
  • Water has molecular mass 18 u.
  • Carbon dioxide has molecular mass 44 u.
  • Nitric acid has molecular mass 63 u.
⚛️

Formula Unit Mass

📋
Table of Contents
12 sections
🗺️ Overview
The formula unit mass of a substance is the sum of the atomic masses of all atoms present in the formula unit of an ionic compound. It is expressed in atomic mass units (u).
Formula unit mass is used for ionic compounds because they do not exist as individual molecules.
📘 Definition
🤔 Did You Know?
Why Do We Use Formula Unit Mass?

Covalent compounds such as water (\(H_2O\)) and carbon dioxide (\(CO_2\)) exist as independent molecules. Therefore, we calculate their molecular masses.

Ionic compounds such as sodium chloride (\(NaCl\)) and magnesium oxide (\(MgO\)) do not exist as separate molecules. Instead, they exist as large ionic lattices made up of positive and negative ions.

Therefore, the term formula unit mass is used for ionic compounds.

🔢 Formula
⚖️ Comparison
Formula Unit vs Molecule
Molecule Formula Unit
Exists independently. Represents the simplest ratio of ions.
Found in covalent compounds. Found in ionic compounds.
Used for molecular mass. Used for formula unit mass.
\(H_2O\), \(CO_2\) \(NaCl\), \(MgO\)
✏️ Example
Solved Example
Formula Unit Mass of Sodium Chloride (NaCl)
Given: Atomic mass of Sodium = 23 u
Atomic mass of Chlorine = 35.5 u
  1. \[ \begin{aligned} \text{Formula Unit Mass} &= 23 + 35.5\\ &=58.5u \end{aligned} \]
Formula unit mass of sodium chloride = 58.5 u.
Formula Unit Mass of Magnesium Oxide (MgO)
Given: Magnesium = 24 u
Oxygen = 16 u
  1. \[24 + 16=40u\]
Formula unit mass of magnesium oxide = 40 u.
Formula Unit Mass of Calcium Carbonate (CaCO₃)
Given: Ca = 40 u
C = 12 u
O = 16 u
  1. \[\begin{aligned}40 + 12 + (3\times16)&=40 + 12 + 48\\&=100u\end{aligned}\]
Formula unit mass of calcium carbonate = 100 u.
🌟 Important Formula Unit Masses
⚖️ Difference Between Molecular Mass and Formula Unit Mass
Molecular Mass Formula Unit Mass
Used for molecules. Used for ionic compounds.
Applies to covalent compounds. Applies to ionic compounds.
\(H_2O\), \(CO_2\) \(NaCl\), \(MgO\)
Molecule exists independently. Formula unit represents the simplest ion ratio.
⚡ Exam Tip
❌ Common Mistakes
  • Using molecular mass for ionic compounds.
  • Ignoring subscripts while calculating masses.
  • Confusing formula units with molecules.
  • Using atomic number instead of atomic mass.
  • Forgetting to multiply oxygen atoms by their subscripts.
📋 Case Study
CBSE Competency-Based Question

A student states that the molecular mass of sodium chloride is 58.5 u.

Is the statement scientifically correct?

Solution

Sodium chloride is an ionic compound and does not exist as independent molecules.

Therefore, the correct term is:

\[ \text{Formula Unit Mass of NaCl} = 58.5u \]

Hence, the student's terminology is incorrect.

⚡ Quick Revision
  • Formula unit mass is used for ionic compounds.
  • It is calculated exactly like molecular mass.
  • NaCl has formula unit mass 58.5 u.
  • MgO has formula unit mass 40 u.
  • CaCO₃ has formula unit mass 100 u.
  • Ionic compounds have formula units, not molecules.
⚛️

Important Points for Quick Revision

📋
Table of Contents
7 sections
🗺️ Overview
This section provides a concise revision of the most important concepts, laws, definitions, formulae, and examination facts from the chapter Atoms and Molecules.
⚖️ Laws
Fundamental Laws of Chemical Combination<
  • Law of Conservation of Mass: During a chemical reaction, mass is neither created nor destroyed. The total mass of reactants is always equal to the total mass of products.
  • Law of Definite Proportions: A pure chemical compound always contains the same elements combined in a fixed proportion by mass, irrespective of its source or method of preparation.
📘 Definition
Important Definitions
🗒️ Important Atomicities to Remember
Important Atomicities to Remember
Element Formula Atomicity
Helium \(He\) 1
Hydrogen \(H_2\) 2
Nitrogen \(N_2\) 2
Oxygen \(O_2\) 2
Ozone \(O_3\) 3
Phosphorus \(P_4\) 4
Sulphur \(S_8\) 8
🌟 Importance
Common Valencies
🔢 Formula
Important Formulae
🔎 Key Fact
High-Yield Examination Facts
· Updated
NCERT Science · Class IX · Chapter 3

Atoms and Molecules

From Dalton’s bold postulates to Avogadro’s giant number — master the building blocks of matter with interactive tools, concept clarity, and deep practice.

⚛ Dalton’s Atomic Theory 🔬 Laws of Chemical Combination 🧮 Mole Concept 📐 Atomic & Molecular Mass 🧪 Chemical Formulae ⚖ Avogadro’s Number
Core Concepts

Eight foundational concept blocks with clear explanations, real-world connections, and key takeaways.

⚖
Concept 01
Laws of Chemical Combination
The two foundational laws that made atomic theory necessary

Law of Conservation of Mass (Lavoisier, 1774): In any chemical reaction, the total mass of all reactants equals the total mass of all products. Mass is neither created nor destroyed. If you burn 12 g of carbon in 32 g of oxygen, you get exactly 44 g of carbon dioxide — no mass disappears into thin air.

Law of Definite Proportions (Proust, 1799): A pure chemical compound always contains the same elements combined in a fixed mass ratio, regardless of its source or method of preparation. Water is always H : O = 1 : 8 by mass, whether from a river, lab synthesis, or mineral extraction.

Conservation of mass means no atoms are created or destroyed in a reaction — they only rearrange.
Definite proportions prove compounds are not random mixtures but have fixed atomic arrangements.
Both laws were direct experimental evidence that matter consists of discrete particles.
Water: H : O = 1 : 8 by mass. CO₂: C : O = 3 : 8 by mass.
🔵
Concept 02
Dalton’s Atomic Theory
The first scientific model — atoms as indivisible, indestructible spheres

John Dalton (1808) proposed the first comprehensive atomic theory to explain the two laws of chemical combination. His postulates were revolutionary for their time, though modern science has partially revised them.

All matter is made of atoms. Atoms are indivisible and indestructible tiny particles.
Atoms of the same element are identical in mass and chemical properties.
Atoms of different elements differ in mass and chemical properties.
Compounds form by combination of atoms of different elements in small whole-number ratios.
Atoms cannot be created or destroyed in chemical reactions — only rearranged.
Modern revision: Atoms ARE divisible (sub-atomic particles exist). Isotopes show same-element atoms can differ in mass.
🔤
Concept 03
Atomic Symbols & Atomic Mass
How we name atoms and measure their incredibly tiny masses

Symbols: Berzelius introduced 1–2 letter symbols. First letter always uppercase, second always lowercase. Some come from Latin: Fe (Ferrum = Iron), Na (Natrium = Sodium), K (Kalium = Potassium), Cu (Cuprum = Copper), Pb (Plumbum = Lead), Ag (Argentum = Silver), Au (Aurum = Gold).

Atomic Mass Unit (u / amu): One u is defined as ¹/₁₂ the mass of one Carbon-12 atom. This gives H ≈ 1 u, O = 16 u, C = 12 u exactly. Using u avoids working with impossibly small numbers like 1.67 × 10⁻²⁴ g.

H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, K=39, Ca=40, Fe=56, Cu=64, Zn=65, Ag=108, Pb=207 (all in u)
Atomic mass is a relative measure — it compares atoms to 1/12 of C-12. It has no unit of its own, though u is often written.
Cl has mass 35.5 u because it is an average of isotopes Cl-35 (75%) and Cl-37 (25%).
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Concept 04
Molecules, Valency & Chemical Formulae
How atoms group, the rules of valency, and reading formulae correctly

Molecules are the smallest particles of elements or compounds capable of independent existence. Elemental molecules: monoatomic (noble gases), diatomic (H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂), triatomic (O₃), polyatomic (P₄, S₈).

Valency is the combining capacity of an atom. H and Na = 1; O, Mg = 2; Al, N = 3; C = 4. To write a formula, use the criss-cross rule: swap the valencies of the combining ions and use them as subscripts, then simplify.

Molecular Mass = sum of atomic masses of all atoms in the molecule. H₂O: 2(1)+16 = 18 u. Glucose C₆H₁₂O₆: 6(12)+12(1)+6(16) = 180 u.

Diatomic elements: HONClBrIF — H₂ O₂ N₂ Cl₂ Br₂ I₂ F₂
Criss-cross: Mg²⁺ + Cl⁻ → MgCl₂ (Mg gets subscript 1, Cl gets subscript 2)
Always simplify ratios after criss-cross: Ca²⁺ + O²⁻ → Ca₂O₂ → CaO (same valency = 1:1 ratio)
Formula unit mass (not molecular mass) is used for ionic compounds like NaCl, MgCl₂.
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Concept 05
Ions & Polyatomic Ions
Charged particles and the special groups that behave as a single unit

Ions are atoms or groups of atoms that carry an overall electric charge. Cations (positive) form when atoms lose electrons — typical of metals. Anions (negative) form when atoms gain electrons — typical of non-metals.

Polyatomic ions are groups of atoms acting as one unit with an overall charge. When writing formulae with polyatomic ions, always enclose the ion in brackets before applying the criss-cross rule if its subscript will be greater than 1.

Common cations: Na⁺, K⁺, H⁺, Ag⁺, Ca²⁺, Mg²⁺, Al³⁺, Fe²⁺, Fe³⁺, Cu²⁺, Zn²⁺, NH₄⁺
Common anions: Cl⁻, Br⁻, OH⁻, NO₃⁻, HCO₃⁻, O²⁻, S²⁻, SO₄²⁻, CO₃²⁻, PO₄³⁻
Al³⁺ + OH⁻ → Al(OH)₃ — brackets are essential here
NH₄⁺ + SO₄²⁻ → (NH₄)₂SO₄
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Concept 06
The Mole Concept
The chemist’s counting unit — a bridge between atoms and grams

Atoms are too small to count individually. Chemists use the mole as a counting unit, just as a dozen = 12. One mole = 6.022 × 10²³ particles (Avogadro’s Number, Nₐ). This number was chosen so that the molar mass in grams equals the atomic/molecular mass in u — a magical and useful coincidence.

Molar mass: The mass of one mole of a substance in grams. Molar mass of O = 16 g/mol; of H₂O = 18 g/mol. One mole of any gas at STP occupies 22.4 L (molar volume).

Nₐ = 6.022 × 10²³ mol⁻¹ (Avogadro’s Number)
Moles = Given mass ÷ Molar mass   (n = m/M)
Number of particles N = n × Nₐ
Mass of 1 atom or molecule = Molar mass ÷ Nₐ
At STP, 1 mole of any gas = 22.4 L
Molar mass numerically equals atomic/molecular mass but unit changes: u → g/mol
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Concept 07
Formula Unit Mass of Ionic Compounds
Why ionic compounds don’t have “molecules” — and how to calculate their mass

Ionic compounds like NaCl do not exist as discrete molecules. They exist as a giant three-dimensional lattice of cations and anions. Therefore we use formula unit mass — the sum of atomic masses of all atoms in one formula unit.

NaCl formula unit mass = 23 + 35.5 = 58.5 u. One mole of NaCl = 58.5 g, containing Nₐ formula units, Nₐ Na⁺ ions, and Nₐ Cl⁻ ions.

NaCl = 58.5 u  |  CaCO₃ = 100 u  |  NaOH = 40 u  |  H₂SO₄ = 98 u  |  HCl = 36.5 u
Use “molecular mass” only for covalent compounds. Use “formula unit mass” for ionic compounds.
For ionic compounds: molar mass (g/mol) = formula unit mass (u) numerically. The calculation method is identical.
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Concept 08
Avogadro’s Hypothesis & Diatomic Molecules
Why Gay-Lussac’s law forced scientists to accept diatomic molecules

Gay-Lussac (1808) discovered gases react in simple whole-number volume ratios. But Dalton’s theory could not explain why 1 vol H₂ + 1 vol Cl₂ gives 2 vol HCl (not 1 vol). Avogadro (1811) resolved this by proposing that equal volumes of gases at same T and P contain equal numbers of molecules — and that elemental gases like H₂ and Cl₂ are diatomic, so each molecule can split and recombine.

Equal volumes of gases at same T and P contain equal numbers of molecules
This is why H₂, O₂, N₂, etc. are diatomic — Dalton incorrectly assumed H and O were monoatomic
1 mol of any gas at STP = 22.4 L is a direct consequence of this hypothesis
Avogadro’s hypothesis was ignored for 50 years until Cannizzaro unified it with atomic theory in 1860
Formula Reference

Every key formula and relationship — logically grouped and exam-ready.

Mole Conversions
Formula 01
Number of Moles (from mass)
n = m / M
n = moles  |  m = given mass (g)  |  M = molar mass (g/mol). M numerically equals atomic/molecular mass in u.
Formula 02
Number of Particles
N = n × Nₐ
N = number of particles  |  n = moles  |  Nₐ = 6.022 × 10²³ mol⁻¹ (Avogadro’s number).
Formula 03
Mass from Number of Particles
m = (N / Nₐ) × M
First find moles from N by dividing by Nₐ, then multiply by molar mass M.
Formula 04
Mass of One Atom or Molecule
m₁ = M / Nₐ
Divide molar mass by Avogadro’s number. Result is in grams. E.g. mass of 1 O atom = 16/(6.022×10²³) g.
Formula 05
Volume at STP (ideal gas)
V = n × 22.4 L
Valid for ideal gases at STP (0°C, 1 atm). 1 mole of any gas = 22.4 L at STP.
Formula 06
Moles from Gas Volume at STP
n = V / 22.4
V in litres. Combine with N = n × Nₐ to find number of molecules in any gas volume.
Mass Calculations
Formula 07
Molecular Mass
M = Σ(atomic mass × count)
H₂O = 2(1)+16 = 18 u  |  CO₂ = 12+2(16) = 44 u  |  NH₃ = 14+3(1) = 17 u
Formula 08
Formula Unit Mass (Ionic)
M = Σ(atomic mass × subscript)
Identical calculation to molecular mass. NaCl = 23+35.5 = 58.5 u  |  CaCO₃ = 40+12+48 = 100 u
Formula 09
Mass Ratio of Elements in Compound
Ratio = (nₐ × Mₐ) : (nₒ × Mₒ)
Water H:O = 2(1) : 1(16) = 1:8 by mass. This ratio is fixed regardless of quantity taken.
Formula 10
Definition of Atomic Mass Unit
1 u = &frac112; mass of C-12 atom
1 u = 1.66 × 10⁻²⁴ g = 1.66 × 10⁻²⁷ kg. All atomic masses are expressed relative to this standard.
The Mole Triangle — Visual Reference
Moles (n) Mass (m, g) Particles (N) × M (g/mol) ÷ M (g/mol) × Nₐ ÷ Nₐ m = (N / Nₐ) × M

Navigate between mass, moles, and number of particles using this triangle

Tips, Tricks & Common Mistakes

Master the exam mindset — know exactly where marks are won and lost.

Smart Strategies
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Memorise the 7 diatomic elements as “HONClBrIF” — they appear constantly in formulae and equation balancing.
✓
Always simplify the formula after criss-crossing. Write CaO not Ca₂O₂; write AlN not Al₃N₃.
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Write Nₐ = 6.022 × 10²³ the moment the exam paper arrives — it is needed in almost every mole problem.
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For “number of atoms” questions: moles × Nₐ × atoms per molecule. Always multiply the final step by the count of atoms in the molecule.
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Use the Mole Triangle: identify which two of {mass, moles, particles} you have, then apply the right operation to reach the third.
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For ionic compounds write “formula unit mass” not “molecular mass” — examiners deduct marks for this distinction.
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Latin-symbol elements to memorise: Fe (Iron), Cu (Copper), Ag (Silver), Au (Gold), Na (Sodium), K (Potassium), Pb (Lead), Hg (Mercury), Sn (Tin).
Common Mistakes
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Writing “molecular mass” for NaCl, MgCl₂, CaCO₃. These are ionic — use formula unit mass.
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Forgetting to simplify after criss-cross. MgO not Mg₂O₂ (both have valency 2 → 1:1 ratio).
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Confusing atomic mass (u) with molar mass (g/mol). Same number, completely different scale and context.
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Not multiplying by subscripts. CO₂: students write 12+16 = 28 instead of 12+2(16) = 44.
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Mixing molecules with atoms. 1 mol H₂O has Nₐ molecules but 3Nₐ atoms (2H + 1O per molecule).
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Omitting brackets for polyatomic ions. Al(OH)₃ written as AlOH₃ completely changes the meaning.
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Stating Dalton’s theory as still fully correct. Modern corrections: atoms ARE divisible; isotopes exist.
Quick Mnemonics & Memory Hooks
HONClBrIF — pronounce “honk-luh-brief”. The 7 diatomic elements: H₂ O₂ N₂ Cl₂ Br₂ I₂ F₂
6.022 × 10²³ — say it aloud 10 times right now. It’s the chemist’s dozen and the most important constant in this chapter.
Units cancel check — g ÷ (g/mol) = mol. If your units cancel to give moles, you’re doing it right.
Latin metals shortlist — “Fe Cu Ag Au Na K Pb Hg Sn” — memorise this row of symbols for objective questions.
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Click elements to see a Bohr model with atomic mass, symbol, valency, protons, electrons, and neutrons.
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NCERT Class 9 Science Chapter 3 Notes: Atoms and Molecules
NCERT Class 9 Science Chapter 3 Notes: Atoms and Molecules — Complete Notes & Solutions · academia-aeternum.com
Atoms and Molecules form the foundation of Chemistry, helping us understand the makeup of everything in our universe. This chapter introduces the concept of atoms—the smallest indivisible particles of matter postulated by ancient philosophers and confirmed through scientific discoveries. You'll learn how scientists like Dalton, Lavoisier, and others developed ideas about atoms, elements, and molecules, and how these concepts explain the formation of substances we see around us. The chapter…
🎓 Class 9 📐 Science 📖 NCERT ✅ Free Access 🏆 CBSE · JEE
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    Frequently Asked Questions

    An atom is the smallest particle of an element that retains its chemical properties.

    A molecule is a group of two or more atoms chemically bonded together.

    Dalton’s theory states matter is made of indivisible atoms, identical for each element, combining in fixed ratios.

    A chemical formula represents a compound using symbols and shows elements’ proportions.

    Atomic mass unit is a standard mass unit equal to 1/12 the mass of one carbon-12 atom.

    Atomicity is the number of atoms in a molecule of an element or compound.

    A mole is the amount of substance containing Avogadro’s number of particles (6.022 × 10²³).

    Mass is neither created nor destroyed during a chemical reaction.

    Elements combine in a fixed mass ratio to form a compound.

    It is the ratio of the average atomic mass of an element to 1/12 of a carbon-12 atom’s mass.

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