Class 9 Science Chapter 3 Laws of Chemical Combination Atoms And Molecules NCERT

Atoms and Molecules: Laws of Chemical Combination, Atomic Mass, and Molecular Formulae

This chapter in CBSE Class 9 Science introduces the fundamental laws that govern chemical reactions, the concept of atoms as building blocks of matter, the mole concept, and the determination of molecular and ionic compound formulae. Building on the earlier understanding of matter, this chapter explains how the atomic theory provides a consistent explanation for the observable patterns in chemical reactions — why substances always combine in fixed proportions, and why the total mass never changes during a chemical reaction.

The law of conservation of mass, first stated by Antoine Lavoisier in 1789, states that mass can neither be created nor destroyed in a chemical reaction — the total mass of the reactants equals the total mass of the products. For example, when 10 g of calcium carbonate is heated strongly, it decomposes into 5.6 g of calcium oxide and 4.4 g of carbon dioxide (10 = 5.6 + 4.4). The law of constant proportions (or definite proportions), established by Joseph Proust, states that any pure chemical compound always contains the same elements in the same proportion by mass, regardless of the source or method of preparation. For example, water (H₂O) always has a mass ratio of 1:8 hydrogen to oxygen — whether from a river or a laboratory. Pure carbon dioxide (CO₂) always contains carbon and oxygen in the mass ratio 3:8. If 3 g of carbon reacts with 8 g of oxygen, 11 g of CO₂ is produced, and if there is any excess of either element, it remains unreacted. The law of multiple proportions (Dalton, 1803) states that if two elements form more than one compound, the masses of one element that combine with a fixed mass of the other element are in a simple whole-number ratio. For example, carbon and oxygen form CO (12:16 = 3:4) and CO₂ (12:32 = 3:8); the ratio of oxygen masses that combine with fixed carbon is 4:8 = 1:2, a simple whole number ratio.

John Dalton's atomic theory (1803) proposed: (1) All matter is composed of indivisible atoms. (2) Atoms of the same element are identical in mass and properties. (3) Atoms of different elements differ in mass. (4) Atoms combine in simple whole-number ratios to form compounds. (5) In a chemical reaction, atoms are rearranged — they are not created, destroyed, or changed into other atoms. While the first postulate was later revised (atoms are divisible into subatomic particles), the theory remains valid in its core principles and explains all the laws of chemical combination. The atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom, approximately 1.66×10⁻²⁷ kg. The relative atomic mass of an element is its average mass compared to 1/12 of a carbon-12 atom. A molecule is the smallest particle of a compound (or element) that can exist independently and retain the chemical properties of that substance. Elements can exist as monatomic molecules (He, Ar), diatomic molecules (H₂, O₂, N₂, F₂, Cl₂), triatomic (O₃, ozone), etc. Compounds are formed when atoms of different elements combine — the molecular formula shows the actual number of atoms of each element in one molecule. The empirical formula shows the simplest whole-number ratio of atoms. The molecular mass of a compound is the sum of the atomic masses of all atoms in the molecule. The mole is the SI unit for amount of substance: one mole of any substance contains exactly 6.022×10²³ particles (Avogadro's number). The mass of one mole of a substance equals its molecular mass in grams (molar mass). Number of moles = given mass / molar mass.

• Law of conservation of mass: total mass remains constant in chemical reactions. Law of constant proportions: a compound always has the same elements in the same proportion by mass.
• Dalton's atomic theory: all matter is atoms; atoms of an element are identical; compounds form by atoms combining in simple whole-number ratios.
• Atomic mass unit: 1 amu = 1/12 mass of a C-12 atom. Relative atomic mass compares an element's average mass to 1 amu.
• Molecular formula shows actual number of atoms; empirical formula shows simplest ratio. Molecular mass sums atomic masses of all atoms.
• One mole = 6.022×10²³ particles (Avogadro's number); molar mass = mass of one mole (g). Number of moles = mass/molar mass.