Chapter Notes: Atoms and Molecules
CBSE Class 9 Science — Chapter 3
Table of Contents
1. Introduction
2. Laws of Chemical Combination
3. John Dalton's Atomic Theory
4. What Is an Atom?
5. Modern Symbols of Elements
6. Atomic Mass
7. How Do Atoms Exist?
8. What Is a Molecule?
9. Atomicity
10. Molecules of Compounds
11. What Is an Ion?
12. Writing Chemical Formulae
13. Formulae of Simple Compounds
14. Molecular Mass
15. Formula Unit Mass
16. Additional Information & Study Resource Details
1. Introduction
Atoms and molecules are the fundamental building blocks of matter. A clear understanding of atoms and molecules explains why substances combine in particular ways, why physical and chemical properties differ from one substance to another, and how new substances are formed in chemical reactions.
Maharishi Kanad and Pakudha Katyayama in ancient India proposed that matter could be divided into smaller, indivisible particles called Parmanu.
Democritus and Leucippus in ancient Greece proposed a similar idea: matter is made up of indivisible particles called atoms.
These early ideas were philosophical and lacked experimental proof until the development of modern chemistry in the 18th century.
In the late 18th century, Antoine L. Lavoisier established quantitative methods in chemistry and laid the foundations of modern chemical science by formulating laws about chemical combinations.
Lavoisier and Joseph L. Proust performed careful experiments that led to two important laws: the Law of Conservation of Mass and the Law of Constant Proportions, also called the Law of Definite Proportions.
These laws guided later work and helped John Dalton formulate his atomic theory, which explained why these laws hold true for chemical reactions.
2. Laws of Chemical Combination
Two fundamental experimental laws describe how substances combine in chemical reactions: the Law of Conservation of Mass and the Law of Constant Proportions. These laws provide the experimental basis for the atomic view of matter.
Law of Conservation of Mass
The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of the reactants is equal to the total mass of the products.
Total mass of reactants = Total mass of products
Example of Law of Conservation of Mass: Careful experiments, such as mixing chemical solutions in closed containers and measuring the mass before and after the reaction, show that the measured total mass remains unchanged. This supports the Law of Conservation of Mass.
Law of Constant Proportion (Law of Definite Proportions)
The Law of Constant Proportion states that a pure chemical compound always contains the same elements combined together in the same fixed proportion by mass, irrespective of its source.
Example: Pure water always contains hydrogen and oxygen in the mass ratio 1 : 8. This proportion is the same whether the water comes from a river, a well, or rain.
3. John Dalton's Atomic Theory
To explain the laws of chemical combination, John Dalton proposed an atomic theory that provided an experimental and conceptual basis for understanding atoms and compounds.
Postulates of Dalton's Atomic Theory
1. All matter is made up of extremely small particles called atoms.
2. Atoms are indivisible by chemical means and remain unchanged during chemical reactions.
3. Atoms of the same element are identical in mass and properties.
4. Atoms of different elements have different masses and properties.
5. Atoms combine in simple whole-number ratios to form compounds.
6. The relative number and types of atoms in a given compound are constant, giving the compound a fixed composition.
Background on John Dalton
John Dalton was born in 1766. His atomic hypothesis explained the Law of Conservation of Mass and the Law of Definite Proportions quantitatively and conceptually.
Modern discoveries have shown that atoms can be divided into smaller subatomic particles and that atoms of the same element may have different masses because of isotopes. However, Dalton's theory remains an important foundation of modern chemistry, especially for understanding chemical reactions.
4. What Is an Atom?
An atom is the smallest particle of an element that retains the chemical properties of that element and cannot be broken down by chemical means.
Atoms are extremely small; a very large number of atoms are required to form visible matter.
A layer only a few million atoms thick may be comparable in thickness to a sheet of paper.
Atomic Radius
The atomic radius is a measure of the size of an atom. It is typically expressed in nanometres:
1 nm = 10⁻⁹ m
5. Modern Symbols of Elements
Element symbols evolved from early pictorial symbols to the simple one-letter or two-letter symbols used today. The International Union of Pure and Applied Chemistry (IUPAC) standardises these symbols.
Historical background: John Dalton first used symbols to represent atoms. Later, Berzelius proposed using one or two letters derived from the element name to represent elements.
Origin of element names: Some names come from places, such as copper from Cyprus, while others come from colours or other properties.
Modern symbols: Most element symbols are derived from their English names. The first letter is always capitalised and the second letter, if present, is lowercase.
First letter and another letter: Examples include chlorine (Cl) and zinc (Zn).
Names from other languages: Some symbols come from Latin, Greek, or other languages. Examples include iron (Fe) from ferrum; sodium (Na) from natrium; and potassium (K) from kalium.
6. Atomic Mass
Atomic mass is the mass of an atom expressed relative to a standard. The standard used internationally is defined as 1/12th of the mass of a carbon-12 atom. The unit for atomic mass is the unified atomic mass unit, represented by u (also called amu).
Atomic mass is the combined mass of the protons, neutrons, and electrons in an atom. In simple calculations, the mass of the electron is often neglected because it is very small compared with the masses of protons and neutrons.
Atomic masses reported in the periodic table are usually average values that reflect the natural isotopic composition of an element.
7. How Do Atoms Exist?
Many atoms do not exist freely under normal conditions. They combine to form molecules or form ions that aggregate into ionic structures.
Visible matter is composed of huge numbers of molecules or ionic units assembled together.
8. What Is a Molecule?
A molecule is the smallest particle of an element or compound that can exist independently and retain the chemical properties of that substance.
Molecules of Elements
1. Monoatomic: Some elements exist as single atoms in their natural gaseous state, for example, Helium (He) and Argon (Ar).
2. Diatomic: Several non-metals exist as molecules made up of two atoms, for example, Hydrogen (H₂), Oxygen (O₂), Nitrogen (N₂), and Chlorine (Cl₂).
3. Polyatomic: Some elements form molecules containing more than two atoms, for example, Phosphorus (P₄) and Sulphur (S₈).
9. Atomicity
Atomicity is the number of atoms present in one molecule.
10. Molecules of Compounds
When atoms of different elements combine, they form molecules of compounds. These molecules have fixed compositions and properties different from those of their constituent elements.
11. What Is an Ion?
An ion is an atom or a group of atoms that carries an electric charge because of the loss or gain of electrons.
A positively charged ion is called a cation, whereas a negatively charged ion is called an anion.
Compounds formed from metals and non-metals often contain ions. Such compounds are called ionic compounds.
Polyatomic ions are groups of atoms bonded together that carry a net charge. Examples include NO₃⁻, SO₄²⁻, and OH⁻.
Ionic compounds do not generally consist of separate molecules. They contain ions arranged in a regular structure. The simplest whole-number ratio of the ions is represented by the compound's formula unit.
12. Writing Chemical Formulae
The chemical formula of a compound shows which elements are present and the number of atoms of each element in the smallest unit of that compound.
To write formulae, one must know the element symbols and the valencies, or ionic charges, of the atoms or ions involved.
Valency indicates how many electrons an atom gains, loses, or shares when it forms a compound. It can also be understood as the number of bonds an atom typically forms (its combining capacity).
Rules for Formula Writing
The total positive charge and total negative charge in a neutral compound must balance.
When writing formulae for compounds of a metal and a non-metal, write the metal first and the non-metal second (for example, CaO and NaCl).
Use the simplest whole-number ratio of atoms or ions that balances the charges.
When polyatomic ions are present in more than one number, enclose the polyatomic ion in brackets and write the number outside the brackets (for example, Mg(OH)₂).
13. Formulae of Simple Compounds
Binary compounds are formed from two elements. Their formulae can be written by using the valencies of the elements or by balancing the charges of the ions.
Carbon tetrachloride (CCl₄): Carbon has valency 4 and chlorine has valency 1. Therefore, one carbon atom combines with four chlorine atoms to form CCl₄.
Magnesium chloride (MgCl₂): Magnesium has valency 2 and chlorine has valency 1. Therefore, one magnesium atom combines with two chlorine atoms to form MgCl₂.
14. Molecular Mass
The molecular mass of a molecule is the sum of the atomic masses of all the atoms present in the molecule. It is expressed in atomic mass units, represented by u.
Example Calculations
Problem: (a) Calculate the molecular mass of water (H₂O).
(b) Calculate the molecular mass of nitric acid (HNO₃).
Solution:
(a) Atomic mass of hydrogen = 1 u, oxygen = 16 u.
M(H₂O) = 2 × 1 + 16 = 18 u.
(b) Atomic mass of hydrogen = 1 u, nitrogen = 14 u, oxygen = 16 u.
M(HNO₃) = 1 × 1 + 1 × 14 + 3 × 16 = 1 + 14 + 48 = 63 u.
15. Formula Unit Mass
Formula unit mass is the sum of the atomic masses of the atoms present in the formula unit of an ionic compound. It is calculated in the same way as molecular mass, but it is applied to ionic formula units.
Example Calculation
Problem: Calculate the formula unit mass of Calcium Chloride (CaCl₂).
Solution:
The atomic mass of calcium is 40 u, and the atomic mass of chlorine is 35.5 u.
M(CaCl₂) = 40 + 2 × 35.5 = 40 + 71 = 111 u.
16. Additional Information & Resource Details
Document Source: EduRev Class 9 Science Chapter 3 — Atoms and Molecules Chapter Notes.
Course Author: Dr. Manju Sen (Teaching Class 9 for 12+ years, merit rank holder, creator of 6 online EduRev courses).
Document Stats: 55.6K Views | Rating: 4.63 / 5 | Last Updated: Oct 30, 2026.
Target Exams & Syllabi: CBSE Class 9 Science, JSS 3, Grade 9, BMAT.
Summary of Material: Comprehensive topic coverage including laws of chemical combination, atomic and molecular masses, formula writing, atomicity, valency, and practice MCQs with detailed explanations.
