3. Atoms and Molecules

🧪 Chapter 3: Atoms and Molecules – Notes

🔹 1. Laws of Chemical Combination

⚖️ Law of Conservation of Mass

• Mass can neither be created nor destroyed in a chemical reaction.

• Total mass of reactants = Total mass of products.

🧮 Law of Constant Proportions

• In a chemical substance, elements are always present in definite proportions by mass.

• Example: Water (H₂O) always contains 2 parts hydrogen and 16 parts oxygen by mass (ratio 1:8).

🔹 2. What is an Atom?

• An atom is the smallest particle of an element that retains its chemical properties.

• Concept first introduced by Maharishi Kanad in India and Democritus in Greece.

• Modern atomic theory by John Dalton.

📘 Dalton’s Atomic Theory – Main Postulates:

1. All matter is made of tiny particles called atoms.

2. Atoms are indivisible and indestructible.

3. All atoms of an element are identical in mass and properties.

4. Compounds form when atoms of different elements combine in fixed ratios.

5. Chemical reactions involve reorganization of atoms.

🔹 3. What is a Molecule?

• Molecule: Two or more atoms chemically bonded together.

• Can be of:

  • Same element: e.g., O₂, N₂ (called diatomic molecules)

  • Different elements: e.g., H₂O, CO₂

Types of Molecules:

• Monoatomic: He, Ne (noble gases)

• Diatomic: H₂, O₂, N₂

• Polyatomic: O₃ (ozone), P₄ (phosphorus), S₈ (sulphur)

🔹 4. Ions

• Ion: A charged particle (atom or group of atoms).

  • Cation: Positively charged ion (loss of electrons)
    e.g., Na⁺, Mg²⁺

  • Anion: Negatively charged ion (gain of electrons)
    e.g., Cl⁻, O²⁻

🔹 5. Chemical Formulae

• Represents the composition of a compound.

• Example:

  • Water → H₂O (2 hydrogen atoms, 1 oxygen atom)

  • Carbon dioxide → CO₂ (1 carbon atom, 2 oxygen atoms)

Rules to Write Chemical Formula:

1. Cation is written first, followed by anion.

2. Cross-multiply the valencies to balance the charges.

3. Reduce to the simplest whole-number ratio if needed.

Examples:

• Sodium chloride: Na⁺ + Cl⁻ → NaCl

• Magnesium chloride: Mg²⁺ + Cl⁻ → MgCl₂

• Aluminium oxide: Al³⁺ + O²⁻ → Al₂O₃

🔹 6. Molecular Mass and Formula Unit Mass

💡 Molecular Mass:

• Sum of atomic masses of all atoms in a molecule.
  Example:

  • H₂O = (2 × 1) + (1 × 16) = 18 u

💡 Formula Unit Mass:

• Used for ionic compounds.

  • NaCl = (23 + 35.5) = 58.5 u

🔹 7. Mole Concept

• 1 mole = 6.022 × 10²³ particles (Avogadro Number)

• Used to count atoms, molecules, ions etc.

🔢 Molar Mass:

• Mass of 1 mole of a substance in grams = Molecular mass in u.

Example:

• H₂O → Molecular mass = 18 u
  So, 1 mole of water = 18 g = 6.022 × 10²³ molecules

🔹 8. Calculations using Mole Concept

• Number of moles =

  Given mass (g)Molar mass (g/mol)\frac{\text{Given mass (g)}}{\text{Molar mass (g/mol)}}Molar mass (g/mol)Given mass (g)​

• Number of particles =

  Moles×6.022×1023\text{Moles} \times 6.022 \times 10^{23}Moles×6.022×1023

🧠 Important Terms

📝 Example Questions

1. What is the molecular mass of CO₂?
   = 12 (C) + 16×2 (O) = 44 u

2. Write the chemical formula of calcium chloride.
   = Ca²⁺ + Cl⁻ → CaCl₂

3. Calculate number of moles in 36g of water.
   = 36 / 18 = 2 moles

4. How many molecules are present in 2 moles of CO₂?
   = 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules