Chapter 3 · Class 9 · Chemistry · CBSE/NCERT
Atoms and Molecules
How can 18 g of water contain 6 × 10²³ molecules? Why does water always have hydrogen and oxygen in the same ratio? This chapter covers laws of chemical combination, Dalton's atomic theory, symbols, valency, formulae, molecular mass, and the mole concept.
AtomsMoleculesValencyChemical FormulaeMolar MassMole Concept
Complete Chapter Roadmap — 12 Topics
1
Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical reaction
Law Statement
In a chemical reaction: Total mass of reactants = Total mass of products. Mass is conserved; it only changes form.
The law of conservation of mass states that in a chemical reaction, mass is neither created nor destroyed. The total mass of reactants equals the total mass of products.
If you burn magnesium ribbon in a closed container and carefully collect all the products (including the magnesium oxide formed and the oxygen consumed), you will find the total mass before the reaction equals the total mass after.
This law was established experimentally by Antoine Lavoisier. It forms the basis of balancing chemical equations — whatever is on the left side of the arrow must equal the right side by mass.
2
Law of Constant Proportions
A compound always has the same elements in the same ratio by mass
Law Statement
A pure chemical compound always has the same elements combined in the same fixed ratio by mass, regardless of source or method of preparation.
The law of constant proportions (also called the law of definite proportions) states that a pure compound always contains the same elements in the same fixed proportion by mass, regardless of the source or method of preparation.
Pure water from the Ganga, from a Mumbai tap, from a bore well in Rajasthan, or prepared in the lab — all contain hydrogen and oxygen in the same mass ratio of 1 : 8.
This is why compounds are different from mixtures. You can make dilute or concentrated nimbu pani (a mixture) by varying proportions. But water is always H₂O in a fixed mass ratio — you cannot make "dilute water" or "concentrated water."
Quick Check
In a sample of water, hydrogen and oxygen are in 1 : 8 ratio. If you have 18 g of water, how much hydrogen is present?
Answer: Ratio is 1 : 8 by mass, total = 9 parts. In 18 g: hydrogen = g.
Answer: Ratio is 1 : 8 by mass, total = 9 parts. In 18 g: hydrogen = g.
3
Dalton's Atomic Theory
The first scientific model of the atom
John Dalton proposed his atomic theory in 1808 to explain the laws of chemical combination.
Main postulates: (1) All matter is made of tiny indivisible particles called atoms. (2) Atoms of the same element are identical in mass and properties. (3) Atoms of different elements have different masses and properties. (4) Atoms cannot be created or destroyed in a chemical reaction. (5) Compounds are formed when atoms of different elements combine in simple whole-number ratios.
Some of Dalton's postulates have been revised since — atoms are not indivisible (they have subatomic particles), and atoms of the same element can have different masses (isotopes). But the theory was a major scientific breakthrough that explained chemical combination successfully.
4
Atoms and Atomic Symbols
Atoms are represented by short international symbols
Important Symbols
H (Hydrogen) | O (Oxygen) | C (Carbon) | N (Nitrogen) | Na (Sodium, from Natrium) | K (Potassium, from Kalium) | Fe (Iron, from Ferrum) | Cu (Copper, from Cuprum) | Ag (Silver, from Argentum) | Au (Gold, from Aurum) | Ca (Calcium) | Mg (Magnesium) | Al (Aluminium) | Cl (Chlorine) | S (Sulphur)
An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are too small to be seen directly; their sizes are on the order of 0.1 nanometres.
Each element has a chemical symbol — a one or two-letter abbreviation. Many symbols come from the element's Latin or Greek name, which is why some symbols do not match the English name.
5
Atomic Mass
Relative mass of an atom on a standard scale
Key Atomic Masses to Remember
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
The actual mass of an atom is extremely small and inconvenient to use in everyday calculations. So atomic mass is defined as a relative mass — the mass of an atom relative to one-twelfth the mass of a carbon-12 atom.
This relative scale is dimensionless and denoted u (unified atomic mass unit) or amu. On this scale, hydrogen ≈ 1 u, carbon = 12 u, oxygen = 16 u, sodium = 23 u, iron = 56 u.
Atomic mass is a weighted average of the masses of all naturally occurring isotopes of the element. This is why atomic mass values are often not whole numbers (e.g., chlorine ≈ 35.5 u).
6
Molecules and Chemical Formulae
Groups of atoms bonded together to form a molecule
A molecule is the smallest particle of a substance that can exist independently and retains all the chemical properties of that substance.
Molecules of elements may contain one type of atom: oxygen exists as O₂, nitrogen as N₂, hydrogen as H₂, and ozone as O₃. Noble gases like helium and argon exist as single atoms.
Molecules of compounds contain atoms of different elements: water (H₂O), carbon dioxide (CO₂), glucose (C₆H₁₂O₆), and common salt is an ionic compound (NaCl). A chemical formula shows the types and numbers of atoms in a molecule.
7
Ions and Valency
Charged particles and the combining capacity of atoms
Common Polyatomic Ions
NH₄⁺ (ammonium) | OH⁻ (hydroxide) | NO₃⁻ (nitrate) | SO₄²⁻ (sulphate) | CO₃²⁻ (carbonate) | HCO₃⁻ (bicarbonate) | PO₄³⁻ (phosphate)
An ion is a charged particle formed when an atom gains or loses electrons. A cation is a positively charged ion (loses electrons): Na⁺, Ca²⁺, Al³⁺. An anion is a negatively charged ion (gains electrons): Cl⁻, O²⁻, N³⁻.
Valency is the combining capacity of an atom — how many electrons it loses, gains, or shares to achieve stability. Sodium has valency 1; magnesium has valency 2; aluminium has valency 3.
Polyatomic ions are groups of atoms that carry a charge and act as a single unit in reactions: sulphate (SO₄²⁻), carbonate (CO₃²⁻), nitrate (NO₃⁻), ammonium (NH₄⁺), hydroxide (OH⁻), phosphate (PO₄³⁻).
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Writing Chemical Formulae
Cross the valencies to write correct formulae
Formula Writing Steps
1. Write cation then anion.
2. Cross the valencies (valency of one = subscript of other).
3. Simplify the ratio if needed.
4. Use brackets for polyatomic ions with subscript > 1.
2. Cross the valencies (valency of one = subscript of other).
3. Simplify the ratio if needed.
4. Use brackets for polyatomic ions with subscript > 1.
To write the formula of a binary ionic compound: (1) write the cation first, then the anion. (2) cross the valencies — the valency of one becomes the subscript of the other. (3) simplify if needed.
Sodium oxide: Na⁺ and O²⁻. Valencies are 1 and 2 respectively. Crossing: Na₂O. Simplify: Na₂O (already in lowest ratio).
Calcium chloride: Ca²⁺ and Cl⁻. Valencies 2 and 1. Crossing: CaCl₂.
For polyatomic ions, use brackets when the subscript is more than 1. Calcium hydroxide: Ca²⁺ and OH⁻. Formula = Ca(OH)₂.
Solved Example
Aluminium sulphate
Al³⁺ and SO₄²⁻. Valencies: 3 and 2. Cross them: Al₂(SO₄)₃.
Check: 2 Al³⁺ carry +6 charge; 3 SO₄²⁻ carry −6 charge. Total charge = 0. ✓
Check: 2 Al³⁺ carry +6 charge; 3 SO₄²⁻ carry −6 charge. Total charge = 0. ✓
Quick Practice
Write the formula of magnesium nitrate.
Answer: Mg²⁺ and NO₃⁻. Valencies: 2 and 1. Formula = Mg(NO₃)₂.
Answer: Mg²⁺ and NO₃⁻. Valencies: 2 and 1. Formula = Mg(NO₃)₂.
9
Molecular Mass and Formula Mass
Sum of atomic masses of all atoms in a formula
Worked Out Examples
CO₂: 12 + 2×16 = 44 u.
H₂SO₄: 2 + 32 + 4×16 = 98 u.
NaCl: 23 + 35.5 = 58.5 u.
CaCO₃: 40 + 12 + 3×16 = 100 u.
H₂SO₄: 2 + 32 + 4×16 = 98 u.
NaCl: 23 + 35.5 = 58.5 u.
CaCO₃: 40 + 12 + 3×16 = 100 u.
Molecular mass is the sum of atomic masses of all atoms present in one molecule of a substance.
Molecular mass of water (H₂O) = 2 × 1 + 16 = 18 u.
For ionic compounds like NaCl (which does not have individual molecules), we calculate the formula mass: Na + Cl = 23 + 35.5 = 58.5 u.
Solved Example
Molecular mass of glucose
Glucose = C₆H₁₂O₆.
Molecular mass = 6×12 + 12×1 + 6×16 = 72 + 12 + 96 = 180 u.
Molecular mass = 6×12 + 12×1 + 6×16 = 72 + 12 + 96 = 180 u.
10
Mole Concept
A counting unit for atoms and molecules
Avogadro's Number
1 mole of any substance contains particles.
A mole is the SI unit for the amount of substance. One mole of any substance contains the same fixed number of particles — approximately 6.022 × 10²³ particles. This number is called Avogadro's number (Nₐ).
Just as "a dozen" means 12 of anything — 12 eggs, 12 mangoes, 12 atoms — "a mole" means 6.022 × 10²³ of anything. One mole of rice grains, one mole of water molecules, or one mole of electrons all contain 6.022 × 10²³ particles.
The mole bridges the microscopic world (individual atoms and molecules) to the macroscopic world (grams that can be weighed in a chemistry laboratory). This is why chemists use moles instead of trying to count atoms directly.
11
Molar Mass and Number of Particles
Connecting grams, moles, and number of particles
Key Relationships
(moles = given mass ÷ molar mass)
(number of particles = moles × Avogadro's number)
The molar mass of a substance is the mass of one mole of that substance, expressed in grams per mole (g/mol). It has the same numerical value as the molecular mass or atomic mass in u.
Molar mass of water = 18 g/mol (molecular mass = 18 u). So 18 g of water is exactly one mole and contains 6.022 × 10²³ water molecules.
To find the number of moles in a given mass, divide by the molar mass. To find the number of particles, multiply the number of moles by Avogadro's number.
Solved Example
Moles in 36 g of water
Molar mass of H₂O = 18 g/mol.
Number of molecules = molecules.
Number of molecules = molecules.
Show a particles calculation from given mass
How many molecules are present in 9 g of water? (Molar mass of H₂O = 18 g/mol)
moles.
Number of molecules = molecules.
Quick Practice
How many moles are present in 44 g of CO₂? (C = 12, O = 16)
Answer: Molar mass of CO₂ = 44 g/mol. mole.
Answer: Molar mass of CO₂ = 44 g/mol. mole.
12
Complete Chapter Summary
Key laws, formulas, and ideas for exam revision
Law of conservation of mass: total mass of reactants = total mass of products.
Law of constant proportions: fixed element ratio by mass in a pure compound.
Valency: combining capacity. Cross valencies to write formulae. Use brackets for polyatomic ions.
Molecular mass = sum of atomic masses. Molar mass has same number in g/mol.
Mole:
. Avogadro's number:
.
Number of particles:
.
Atoms and Molecules: revision sheet
Close the notes and explain each idea in your own words. Then check the formulas and attempt the examples before opening their solutions.
Recall checklist
- Chapter Intro
- Laws of Chemical Combination
- Atoms, Symbols, and Molecules
- Valency and Chemical Formulae
- Atomic Mass, Molecular Mass, and Mole
- Easy Problem-Solving Approach
For each topic: state the rule, explain when it applies, give an example and identify one mistake to avoid. Return to any topic you cannot explain without looking.
Formula recap (2)
Atomic Mass, Molecular Mass, and Mole
Atomic Mass, Molecular Mass, and Mole
Number of moles equals given mass divided by molar mass.
Explain your method
Write the given information, choose a method, show the steps, then check units and whether the answer makes sense.
State the law of conservation of mass.
Check worked solution 1
In a chemical reaction, mass is neither created nor destroyed, so the total mass of reactants equals the total mass of products.
Why is water always found with hydrogen and oxygen in a fixed ratio by mass?
Check worked solution 2
Because compounds obey the law of constant proportions, so water always has the same composition by mass.
Apply the concept, avoid the trap
Common mistake: Counting the coefficient of a chemical formula as if it changed only the first element.
Use this reasoning: A coefficient multiplies the entire formula. A subscript applies to the element or bracketed group immediately before it.
Worked application
Count the atoms represented by 3Ca(OH)₂, then find the relative formula mass of Ca(OH)₂ using Ca = 40, O = 16 and H = 1.
Reveal the method and checks
- One formula unit contains 1 Ca, 2 O and 2 H atoms. The bracket subscript 2 multiplies both O and H.
- Three formula units contain 3 Ca, 6 O and 6 H atoms: 15 atoms in total.
- For one formula unit, relative formula mass = 40 + 2(16 + 1) = 74. The leading coefficient 3 changes the number of formula units, not the formula mass of one unit.
Check your understanding
How many oxygen atoms are represented by 2Al₂(SO₄)₃?
Check your reasoning
Each formula unit has 3 × 4 = 12 oxygen atoms. Two formula units contain 24 oxygen atoms. The coefficient also multiplies the aluminium and sulphur counts.
Turn mistakes into revision
Record the question, the step that went wrong and the corrected reason. Try it again tomorrow without the solution, then revisit it after a week. A correct guess still needs an explanation.