Everything around us — air, water, metals, plastic, our own bodies — is made of matter, and all matter is composed of incredibly tiny particles called atoms. An atom is so small that a single full stop printed on this page contains about 10 million carbon atoms arranged side by side. And yet, the entire science of chemistry — reactions, compounds, elements, acids, bases, drugs, materials — is built upon understanding these microscopic particles and how they combine to form molecules.
The concept of the atom dates back to ancient Greek philosophy (Leucippus and Democritus, ~400 BCE), but was first made scientifically rigorous by John Dalton in 1808, whose Atomic Theory explained the quantitative laws of chemical combination. Today, with electron microscopes, we can even “see” individual atoms. For CBSE Class 9 Chemistry Chapter 3 (Atoms and Molecules), this is a foundational chapter that introduces the quantitative language of chemistry: atomic masses, molecular masses, chemical formulae, valency, ions, and the immensely powerful mole concept — the cornerstone of all quantitative chemistry.
- 1. Laws of Chemical Combination
- 2. Dalton’s Atomic Theory
- 3. What is an Atom? Size, Symbols, Atomicity
- 4. Atomic Mass and Atomic Mass Unit (u)
- 5. What is a Molecule? Types and Molecular Mass
- 6. Ions — Cations and Anions
- 7. Writing Chemical Formulae — Valency and Criss-Cross Rule
- 8. The Mole Concept and Avogadro’s Number
- 9. Solved Numerical Problems
- 10. Frequently Asked Questions (FAQ)
1. Laws of Chemical Combination
Before Dalton proposed the Atomic Theory, two fundamental laws were established experimentally to describe how elements combine in chemical reactions:
Proposed by: Antoine Lavoisier (1774)
Example: When hydrogen burns in oxygen to form water:
2H₂ + O₂ → 2H₂O
4 g H₂ + 32 g O₂ = 36 g H₂O
Total mass before = Total mass after = 36 g
Experimental verification: Landolt (1908) weighed sealed glass tubes before and after reactions and found no change in total mass.
Proposed by: Joseph Louis Proust (1799)
Example: Water (H₂O) always contains hydrogen and oxygen in the mass ratio 1:8, regardless of whether it comes from rain, a river, the ocean, or a laboratory synthesis.
H : O = 2 : 16 = 1 : 8 (by mass)
Why this is remarkable: It means a compound is not just a mechanical mixture — it has a fixed chemical identity with a definite formula.
Example: Carbon and oxygen form two compounds: CO (carbon monoxide) and CO₂ (carbon dioxide).
In CO: 12 g carbon combines with 16 g oxygen.
In CO₂: 12 g carbon combines with 32 g oxygen.
Ratio of oxygen masses for same carbon: 16 : 32 = 1 : 2 (simple whole number ratio).
2. Dalton’s Atomic Theory (1808)
John Dalton, an English chemist, proposed the first modern scientific theory of the atom in 1808 to explain the laws of chemical combination. His theory had five key postulates:
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All matter is made of atomsAll matter is composed of extremely small, indivisible, and indestructible particles called atoms. Atoms cannot be created or destroyed in chemical reactions.✘ Limitation: Atoms CAN be divided (into protons, neutrons, electrons) by nuclear reactions.
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Atoms of the same element are identicalAll atoms of a given element are identical in mass, size, and properties. Atoms of different elements differ in mass, size, and properties.✘ Limitation: Isotopes of the same element have different masses (e.g., ¹²C and ¹⁴C, both carbon).
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Compounds are formed by combination of atomsAtoms of different elements combine in simple whole-number ratios to form compounds. The compound has a definite fixed composition.No major limitations — this is fundamentally correct.
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Atoms are rearranged in chemical reactionsA chemical reaction involves the rearrangement of atoms — atoms are neither created nor destroyed. This explains the Law of Conservation of Mass.No major limitations — this is fundamentally correct.
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Atoms of different elements have different massesEach element has a characteristic atomic mass that distinguishes it from other elements. This was the basis of Dalton’s early atomic mass table.No major limitations — fundamentally correct (relative atomic masses are real).
3. What is an Atom? Size, Symbols, Atomicity
Size of an atom: The radius of a hydrogen atom is approximately 1 Å (angstrom) = 10²⁹ m = 0.1 nm. To put this in perspective:
• 1 cm contains about 10 million (10&sup7;) hydrogen atoms in a row.
• A single human hair is about 1 million times wider than a hydrogen atom.
• The smallest visible speck (0.1 mm) contains about 10²² (10 billion trillion) atoms.
Chemical symbols: Each element has a one or two-letter symbol (mostly from its Latin or English name):
• H (Hydrogen), He (Helium), C (Carbon), N (Nitrogen), O (Oxygen), Na (Natrium = Sodium), K (Kalium = Potassium), Fe (Ferrum = Iron), Cu (Cuprum = Copper), Au (Aurum = Gold), Ag (Argentum = Silver), Pb (Plumbum = Lead), Hg (Hydrargyrum = Mercury).
Atomicity: The number of atoms in one molecule of an element. Examples:
• Monoatomic: All noble gases — He, Ne, Ar, Kr, Xe (atomicity = 1)
• Diatomic: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂ (atomicity = 2)
• Triatomic: O₃ (ozone), H₂O, CO₂ (atomicity = 3)
• Tetraatomic: P₄ (white phosphorus) (atomicity = 4)
• Polyatomic: S₈ (sulfur, atomicity = 8)
4. Atomic Mass and Atomic Mass Unit (u)
Reference standard: Carbon-12 isotope (¹²C) is assigned an exact mass of 12 u.
Atomic mass of any element = (mass of one atom of that element) / (1/12 × mass of one carbon-12 atom)
| Element | Symbol | Atomic Mass (u) | Element | Symbol | Atomic Mass (u) |
|---|---|---|---|---|---|
| Hydrogen | H | 1 | Sodium | Na | 23 |
| Carbon | C | 12 | Magnesium | Mg | 24 |
| Nitrogen | N | 14 | Aluminium | Al | 27 |
| Oxygen | O | 16 | Sulphur | S | 32 |
| Fluorine | F | 19 | Chlorine | Cl | 35.5 |
| Neon | Ne | 20 | Potassium | K | 39 |
| Silicon | Si | 28 | Calcium | Ca | 40 |
| Phosphorus | P | 31 | Iron | Fe | 56 |
| Zinc | Zn | 65 | Copper | Cu | 63.5 |
| Silver | Ag | 108 | Gold | Au | 197 |
5. What is a Molecule? Types and Molecular Mass
Examples:
• H₂ (hydrogen gas): 2 H atoms
• O₂ (oxygen gas): 2 O atoms
• O₃ (ozone): 3 O atoms
• N₂ (nitrogen gas): 2 N atoms
• Cl₂ (chlorine gas): 2 Cl atoms
• P₄ (white phosphorus): 4 P atoms
• S₈ (sulfur ring): 8 S atoms
Molecular mass of O₂ = 2 × 16 = 32 u
Molecular mass of O₃ = 3 × 16 = 48 u
Examples & Molecular Masses:
• H₂O (water): 2(1) + 16 = 18 u
• CO₂ (carbon dioxide): 12 + 2(16) = 44 u
• NH₃ (ammonia): 14 + 3(1) = 17 u
• HCl (hydrochloric acid): 1 + 35.5 = 36.5 u
• H₂SO₄ (sulphuric acid): 2(1)+32+4(16) = 98 u
• NaCl (common salt): 23+35.5 = 58.5 u
• CaCO₃ (calcium carbonate): 40+12+3(16) = 100 u
6. Ions — Cations and Anions
Cation (positive ion): Formed when an atom loses electrons. Has fewer electrons than protons.
Examples: Na⁺ (sodium ion), Ca²⁺ (calcium ion), Al³⁺ (aluminium ion), NH₄⁺ (ammonium ion)
Anion (negative ion): Formed when an atom gains electrons. Has more electrons than protons.
Examples: Cl¯ (chloride), O²¯ (oxide), SO₄²¯ (sulphate), NO₃¯ (nitrate), OH¯ (hydroxide), CO₃²¯ (carbonate)
Valency: The combining capacity of an atom is called valency. It equals the number of electrons an atom loses, gains, or shares to achieve a noble gas configuration (full outer shell = 8 electrons for most, 2 for hydrogen).
Common valencies:
• Valency 1: H, Na, K, Li, F, Cl, Br, I, Ag, NH₄⁺
• Valency 2: O, Mg, Ca, Ba, Zn, Fe²⁺, Cu²⁺, SO₄²¯, CO₃²¯
• Valency 3: Al, Fe³⁺, N, PO₄³¯
• Valency 4: C, Si
7. Writing Chemical Formulae — Valency and Criss-Cross Rule
(1) Write the symbols of the combining elements side by side (cation first, anion second for ionic compounds).
(2) Write the valency of each element below its symbol.
(3) Criss-cross the valencies (swap them and write them as subscripts for each other).
(4) Simplify the subscripts to the lowest ratio if needed.
Example: Sodium Chloride (NaCl)
Na (valency 1) | Cl (valency 1)
Criss-cross: Na¹Cl¹ → subscripts become 1 and 1 → simplified = NaCl
Example: Magnesium Oxide (MgO)
Mg (valency 2) | O (valency 2)
Criss-cross: Mg²O² → simplified (divide by 2) = MgO
Example: Aluminium Oxide (Al₂O₃)
Al (valency 3) | O (valency 2)
Criss-cross: Al²O³ → already simplified = Al₂O₃
Note: If the valency of an ion is a polyatomic group (like SO₄²¯), enclose it in parentheses before adding the subscript. E.g., Calcium Sulphate: Ca(SO₄) → subscript 2 on Ca, subscript 1 on SO₄ = CaSO₄.
Mol. mass: 2(27) + 3(32+64) = 54 + 288 = 342 u
Mol. mass: 40 + 2(16+1) = 40 + 34 = 74 u
Mol. mass: 2(56) + 3(96) = 112 + 288 = 400 u
Mol. mass: (14+4) + (14+48) = 18 + 62 = 80 u
8. The Mole Concept and Avogadro’s Number
Moles → Particles: N = n×Nₐ
Moles → Mass: m = n×M
Particles → Moles: n = N/Nₐ
All roads lead through the mole!
9. Solved Numerical Problems
= 2×1 + 1×32 + 4×16
= 2 + 32 + 64
Step 2 — Atoms: N = n × Nₐ = (1/3) × 6.022 × 10²³
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10. Frequently Asked Questions (FAQ)
Avogadro’s number (Nₐ) = 6.022 × 10²³ per mole. It is the number of atoms, molecules, ions, or any other elementary entities present in one mole of a substance.
Why it is important: Avogadro’s number bridges the submicroscopic world (individual atoms and molecules) and the macroscopic world (grams of substance we can weigh in the lab). Without this number, it would be impossible to relate the mass of a substance we weigh on a balance to the number of atoms or molecules involved in a reaction.
Example: If you want to react exactly 1 atom of hydrogen with 1 atom of chlorine to make 1 molecule of HCl, you cannot weigh individual atoms. But you CAN weigh 1 g of H (= 1 mole = 6.022×10²³ atoms) and 35.5 g of Cl (= 1 mole = 6.022×10²³ atoms), and they react perfectly in a 1:1 ratio.
Named after: Amedeo Avogadro (1776–1856), Italian scientist who first proposed that equal volumes of all gases (at the same T and P) contain equal numbers of molecules.
Mole concept: The mole is the SI unit for the amount of substance. One mole of any substance contains exactly 6.022 × 10²³ (Avogadro’s number) elementary entities (atoms, molecules, ions, etc.).
Molar mass: The mass of one mole of a substance in grams, numerically equal to the atomic or molecular mass in u. E.g., H₂O has molecular mass 18 u, so its molar mass = 18 g/mol.
Key formulas:
n (moles) = m (grams) / M (g/mol) → n = m/M
N (particles) = n × Nₐ = n × 6.022×10²³
m = n × M
Practical use: To find how many grams of oxygen react with 4 g of hydrogen to form water:
H₂ + ½O₂ → H₂O (1 mol H₂ reacts with 0.5 mol O₂)
4 g H₂ = 4/2 = 2 mol H₂ → needs 1 mol O₂ = 32 g O₂. The mole makes this calculation straightforward.
Atom: The smallest particle of an element that can take part in a chemical reaction and retains the chemical properties of the element. Atoms cannot be broken down by chemical means (only by nuclear reactions). Most atoms cannot exist independently for extended periods — they either form molecules or ionic lattices. Noble gas atoms (He, Ne, Ar) are stable as single atoms.
Molecule: A group of two or more atoms bonded together by covalent bonds (shared electrons). A molecule is the smallest particle of a substance (element or compound) that can exist independently and has all the chemical properties of that substance.
Key differences:
• An atom is a single unit of an element; a molecule may contain atoms of the same element (e.g., O₂) or different elements (e.g., H₂O).
• Atoms (except noble gases) usually cannot exist independently; molecules can exist independently.
• All molecules are made of atoms; not all atoms form molecules (e.g., noble gases are monatomic).
Examples: H is an atom; H₂ is a molecule. O is an atom; H₂O (water) is a molecule. Na is an atom; NaCl (sodium chloride) is an ionic compound (not a covalent molecule).
The two main laws of chemical combination are:
1. Law of Conservation of Mass (Lavoisier, 1774): In any chemical reaction, the total mass of the reactants equals the total mass of the products. Mass is neither created nor destroyed. For example, when 2H₂ + O₂ → 2H₂O: 4g + 32g = 36g (both sides equal 36g).
2. Law of Definite Proportions (Proust, 1799): A pure chemical compound always contains the same elements combined in the same definite proportion by mass, regardless of the source of the compound. For example, water always contains hydrogen and oxygen in the mass ratio 1:8.
Additional law (Dalton, 1803):
3. Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers. E.g., CO has C:O = 12:16 and CO₂ has C:O = 12:32. Oxygen ratio for same carbon: 16:32 = 1:2.
The criss-cross method is a simple technique for writing the chemical formula of an ionic compound using the valencies (combining capacities) of the ions:
Steps:
(1) Write the symbols of the combining elements/ions side by side (cation first, anion second).
(2) Write the valency of each ion/atom below/above its symbol.
(3) Criss-cross the valencies: the valency of the first element becomes the subscript of the second element, and vice versa.
(4) Reduce the subscripts to the lowest whole number ratio if both are divisible by a common factor.
Examples:
• MgCl₂: Mg (val=2), Cl (val=1) → Mg¹Cl² → MgCl₂
• Al₂O₃: Al (val=3), O (val=2) → Al²O³ → Al₂O₃
• CaSO₄: Ca (val=2), SO₄ (val=2) → Ca²(SO₄)² → simplified → CaSO₄
• Na₂CO₃: Na (val=1), CO₃ (val=2) → Na²(CO₃)¹ → Na₂CO₃
Important rule: If both subscripts are the same, divide by that number to simplify (e.g., Ca²O² → CaO). If the anion is polyatomic (e.g., SO₄, CO₃, NO₃, OH), use parentheses before adding the subscript when the subscript is 2 or more (e.g., Ca(OH)₂, NOT CaOH₂).
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