Page 1: Introduction

All matter is made of very small particles called atoms.

Atoms combine to form molecules.

This chapter explains laws of chemical combination, atomic and molecular masses, mole concept, and calculations.

Page 2: Laws of Chemical Combination

• Law of Conservation of Mass (Lavoisier): Mass neither created nor destroyed.
• Law of Constant Proportions (Proust): Fixed ratio by mass.
• Law of Multiple Proportions (Dalton): Fixed ratio for same elements in different compounds.
• Gay Lussac’s Law of Gaseous Volumes: Gases combine in simple volume ratio.
• Avogadro’s Law: Equal volumes of gases contain equal number of molecules at same T and P.

Page 3: Dalton’s Atomic Theory

Limitations known now (isotopes, subatomic particles).

Page 4: Atom, Molecule, Ion

• Atom: Smallest particle that takes part in reaction.
• Molecule: Group of atoms (same or different) — smallest unit of compound/element.
• Ion: Charged particle (cation +, anion -).

Atomicity: monoatomic (He), diatomic (O₂), triatomic (O₃), polyatomic (P₄).

Page 5: Atomic Mass Unit

1 u = 1/12 mass of carbon-12 atom ≈ 1.66 × 10⁻²⁷ kg.

H = 1 u, C = 12 u, O = 16 u, N = 14 u, Na = 23 u, Cl = 35.5 u.

Page 6: Molecular Mass

Sum of atomic masses of all atoms in molecule.

Page 7: Formula Mass

For ionic compounds (no molecule).

Page 8: Mole Concept

1 mole = 6.022 × 10²³ particles (Avogadro’s number N_A).

1 mole of any substance = its gram atomic/gram molecular mass.

Molar mass = atomic/molecular mass in grams.

Page 9: Calculations Using Mole

• Number of moles n = mass / molar mass
• Number of particles = n × N_A
• Mass = n × molar mass

Page 10: Percentage Composition

% of element = (mass of element in compound / molar mass) × 100

Page 11: Key Formulas Summary

• Molecular mass = Σ atomic masses
• n = m / M
• Particles = n × 6.022 × 10²³
• % composition formula

Page 12: Practice Questions - Easy (1-10)

1. Molecular mass of CO₂.
2. 1 mole particles?
3. Atomic mass of oxygen.
4. Molar mass of water.
5. Mass of 2 moles NaCl.
6. Number of atoms in 0.5 mole O₂.
7. % carbon in CO₂.
8. Formula mass CaCO₃.
9. Avogadro’s number value.
10. Law of conservation example.

Page 13: Practice Questions - Medium (11-20)

11. Moles in 44 g CO₂.
12. Atoms in 18 g H₂O.
13. % hydrogen in H₂O.
14. Find molar mass Al₂(SO₄)₃.
15. Mass of 3.011 × 10²³ molecules N₂.
16. Number of moles in 112 g N₂.
17. % composition of Na₂CO₃.
18. Molecules in 64 g SO₂.
19. Atoms in 4 g He.
20. Empirical formula hint.

Page 14: Practice Questions - Hard (21-30)

21. Find mass of 10²³ atoms C.
22. Volume at STP relation hint.
23. Combined calculations.
24. Percentage composition detailed.
25. Large number conversions.
26. Find number of ions.
27. Multiple step mole problems.
28. Real-life application.
29. Verify law numerically.
30. Advanced Dalton theory.

Page 15: NCERT Exercise Types

Mole calculations, percentage composition, mass-number conversions.

Page 16: Important Constants

N_A = 6.022 × 10²³
Common atomic masses table.

Page 17: Common Mistakes

• Forgetting ×2 for diatomic
• Wrong molar mass
• Confusing atoms vs molecules
• Percentage calculation error
• Scientific notation mistake

Page 18: Previous Year Questions

Mole-mass-particle conversions, percentage composition.

Page 19: Exam Tips

• Write formula first
• Show units
• Use exact N_A value if needed
• Memorise common masses
• Step-by-step calculation

Page 20: Quick Revision Sheet

All formulas, laws, constants.

Page 21: Laws Table

Laws of chemical combination summary.

Page 22: Final Motivation

Chapter 3 complete! Mole concept is super important for Class 10 Chemistry.

Practice calculations daily.

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Page 23: Atomic Mass Table

Common elements list.

Page 24: Calculation Steps

Standard method for mole problems.

Page 25: Extra Examples

More solved numericals.

Page 26: Dalton Theory Points

Detailed postulates.