Class 9 Science Notes Chapter 3 (Atoms and Molecules) – Science Book

Alright class, let's focus on Chapter 3: Atoms and Molecules. This is a fundamental chapter in chemistry, and understanding these concepts is crucial not just for your Class 9 exams but also forms the bedrock for many questions in various government recruitment exams. Pay close attention.

Chapter 3: Atoms and Molecules - Detailed Notes for Exam Preparation

1. Laws of Chemical Combination

These laws were established after extensive experimentation by scientists like Lavoisier and Proust and laid the foundation for understanding how elements combine.

• Law of Conservation of Mass:

  • Stated by: Antoine Lavoisier (often called the 'Father of Modern Chemistry').
  • Statement: Mass can neither be created nor destroyed in a chemical reaction.
  • Explanation: In any chemical reaction, the total mass of the reactants (substances that react) is always equal to the total mass of the products (substances formed).
  • Example: If 10g of calcium carbonate (CaCO₃) is heated, it decomposes into 5.6g of calcium oxide (CaO) and 4.4g of carbon dioxide (CO₂). Total mass of reactants (10g) = Total mass of products (5.6g + 4.4g = 10g).
  • Exam Relevance: Often tested conceptually or through simple calculation problems verifying the law.

• Law of Constant Proportions (or Law of Definite Proportions):

  • Stated by: Joseph Proust.
  • Statement: In a chemical substance, the elements are always present in definite proportions by mass, irrespective of the source or method of preparation.
  • Explanation: A pure chemical compound always contains the same elements combined together in the same fixed ratio by mass.
  • Example: Water (H₂O) always consists of Hydrogen and Oxygen combined in a ratio of 1:8 by mass (Atomic mass H=1u, O=16u; H₂O has 2x1u of H and 1x16u of O, so ratio is 2:16 or 1:8). Whether you take tap water, river water, or water synthesized in a lab, this ratio remains constant.
  • Exam Relevance: Understanding the definition and its implication for the purity and identity of compounds.

2. Dalton's Atomic Theory

John Dalton proposed a theory in 1808 to explain the laws of chemical combination.

• Postulates:

  1. All matter is made up of very tiny particles called atoms, which participate in chemical reactions.
  2. Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction (supports Law of Conservation of Mass).
  3. Atoms of a given element are identical in mass and chemical properties.
  4. Atoms of different elements have different masses and chemical properties.
  5. Atoms combine in the ratio of small whole numbers to form compounds (supports Law of Constant Proportions).
  6. The relative number and kinds of atoms are constant in a given compound.

• Limitations (Drawbacks discovered later):

  • Atoms are not indivisible; they are composed of subatomic particles (protons, neutrons, electrons).
  • Atoms of the same element can have slightly different masses (Isotopes, e.g., ¹²C, ¹³C, ¹⁴C or ³⁵Cl, ³⁷Cl).
  • Atoms of different elements can have the same mass (Isobars, e.g., ⁴⁰Ar, ⁴⁰K, ⁴⁰Ca).
  • The ratio in which atoms combine may not always be simple whole numbers (e.g., in complex organic compounds like sucrose C₁₂H₂₂O₁₁).

• Exam Relevance: Know the postulates and limitations. Dalton's theory was a major step, even with its limitations.

3. Atoms

• Definition: The smallest particle of an element that may or may not exist independently but takes part in a chemical reaction.
• Size: Atoms are extremely small, measured in nanometres (nm). 1 nm = 10⁻⁹ m. (e.g., radius of a hydrogen atom ≈ 10⁻¹⁰ m).
• Symbols:
  • Modern symbols proposed by Berzelius.
  • Approved by IUPAC (International Union of Pure and Applied Chemistry).
  • Usually the first letter (capitalized) or the first two letters (first capitalized, second small) of the element's English name (e.g., H for Hydrogen, He for Helium, Al for Aluminium).
  • Some symbols derived from Latin/German/Greek names (e.g., Na for Sodium - Natrium, K for Potassium - Kalium, Fe for Iron - Ferrum, Au for Gold - Aurum, Ag for Silver - Argentum).
  • Exam Relevance: Be familiar with symbols of common elements (at least first 20, plus common metals like Fe, Cu, Zn, Ag, Au, Pb).

4. Atomic Mass

• Concept: Since atoms are tiny, their absolute mass is very small. We use relative atomic mass.
• Standard: Carbon-12 isotope is assigned a relative atomic mass of exactly 12 atomic mass units.
• Atomic Mass Unit (amu or u): Defined as exactly 1/12th the mass of one atom of Carbon-12. (1 u ≈ 1.66 × 10⁻²⁴ g).
• Atomic Mass of an Element: The average relative mass of its atoms compared to 1/12th the mass of a Carbon-12 atom. (Average considers isotopes).
• Exam Relevance: Know the definition and atomic masses of common elements (H=1u, C=12u, N=14u, O=16u, Na=23u, Mg=24u, S=32u, Cl=35.5u, Ca=40u).

5. Molecules

• Definition: The smallest particle of an element or a compound that is capable of independent existence and shows all the properties of that substance. It is a group of two or more atoms chemically bonded together.
• Molecules of Elements: Formed by the combination of atoms of the same element.
  • Atomicity: The number of atoms constituting a molecule of an element.
    • Monoatomic: Exist as single atoms (e.g., Noble gases - He, Ne, Ar). Atomicity = 1.
    • Diatomic: Exist as molecules of two atoms (e.g., H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂). Atomicity = 2.
    • Triatomic: (e.g., Ozone - O₃). Atomicity = 3.
    • Polyatomic: Contain more than two atoms (e.g., Phosphorus - P₄, Sulphur - S₈). Atomicity = 4 and 8 respectively.
• Molecules of Compounds: Formed by the combination of atoms of different elements in a fixed ratio. (e.g., H₂O, CO₂, NH₃, CH₄).

6. Ions

• Definition: Charged species formed when an atom or molecule loses or gains electrons.
• Cations: Positively charged ions formed by the loss of one or more electrons (e.g., Na⁺, Mg²⁺, Al³⁺, NH₄⁺). Metals typically form cations.
• Anions: Negatively charged ions formed by the gain of one or more electrons (e.g., Cl⁻, O²⁻, N³⁻, SO₄²⁻, OH⁻). Non-metals typically form anions.
• Polyatomic Ions: A group of atoms carrying a net charge, acting as a single unit (e.g., Ammonium NH₄⁺, Hydroxide OH⁻, Nitrate NO₃⁻, Carbonate CO₃²⁻, Sulphate SO₄²⁻, Phosphate PO₄³⁻).
• Valency: The combining capacity of an element or ion. For simple ions, it's equal to the charge on the ion.
• Exam Relevance: Memorize common simple and polyatomic ions along with their charges – essential for writing chemical formulae.

Table of Common Ions:

(Note: Roman numerals indicate the charge for metals with variable valency like Copper and Iron)

7. Writing Chemical Formulae

• Rules:

  1. Write the symbols of the cation (positive ion) first, followed by the anion (negative ion).
  2. Write the valency/charge of each ion below its symbol.
  3. Criss-cross the valencies/charges (use the numerical value only, ignore the sign). These numbers become the subscripts for the ions.
  4. If a subscript is 1, it is not written.
  5. If a polyatomic ion needs a subscript greater than 1, enclose the ion in parentheses before writing the subscript.
  6. Simplify the formula to the simplest whole number ratio if possible (e.g., for Calcium Oxide Ca²⁺ O²⁻, the formula is CaO, not Ca₂O₂).

• Examples:

  • Sodium Chloride: Na⁺ Cl⁻ -> Na₁Cl₁ -> NaCl
  • Magnesium Hydroxide: Mg²⁺ OH⁻ -> Mg₁(OH)₂ -> Mg(OH)₂
  • Aluminium Sulphate: Al³⁺ SO₄²⁻ -> Al₂(SO₄)₃
  • Calcium Carbonate: Ca²⁺ CO₃²⁻ -> Ca₂(CO₃)₂ -> CaCO₃ (Simplified)
  • Ammonium Phosphate: NH₄⁺ PO₄³⁻ -> (NH₄)₃(PO₄)₁ -> (NH₄)₃PO₄

• Exam Relevance: Very important skill. Expect direct questions on writing formulae or identifying correct formulae.

8. Molecular Mass and Formula Unit Mass

• Molecular Mass: The sum of the atomic masses of all the atoms in a molecule of a substance. Expressed in atomic mass units (u).
  • Example: Molecular mass of H₂O = (2 × Atomic mass of H) + (1 × Atomic mass of O) = (2 × 1u) + (1 × 16u) = 18u.
  • Example: Molecular mass of HNO₃ = (1×1u) + (1×14u) + (3×16u) = 1 + 14 + 48 = 63u.
• Formula Unit Mass: The sum of the atomic masses of all atoms in a formula unit of an ionic compound. Calculated the same way as molecular mass, but used specifically for ionic compounds which exist as crystal lattices, not discrete molecules. Expressed in atomic mass units (u).
  • Example: Formula unit mass of NaCl = (1 × At. mass of Na) + (1 × At. mass of Cl) = (1 × 23u) + (1 × 35.5u) = 58.5u.
  • Example: Formula unit mass of CaCl₂ = (1 × 40u) + (2 × 35.5u) = 40u + 71u = 111u.

9. Mole Concept

This is a cornerstone of chemical calculations. It relates mass, number of particles, and volume (for gases).

• Mole: One mole of any species (atoms, molecules, ions, or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.

• Avogadro's Constant (N<0xE2><0x82><0x90>): The number of particles (atoms, molecules, ions) present in exactly one mole of any substance. Its value is fixed: 6.022 × 10²³ particles/mol. This number is called Avogadro's number or Avogadro's constant.

• Molar Mass (M): The mass of one mole of a substance. It is numerically equal to the atomic/molecular/formula unit mass but expressed in grams per mole (g/mol).

  • Molar mass of O atoms = 16 g/mol
  • Molar mass of O₂ molecules = 32 g/mol
  • Molar mass of H₂O molecules = 18 g/mol
  • Molar mass of NaCl = 58.5 g/mol

• Key Relationships (Essential for problem-solving):

  • 1 mole = 6.022 × 10²³ particles (atoms/molecules/ions)
  • 1 mole = Molar Mass (in grams)
  • Number of moles (n) = Given Mass (m) / Molar Mass (M)
  • Number of moles (n) = Given Number of Particles (N) / Avogadro's Number (N<0xE2><0x82><0x90>)
  • Mass (m) = Number of moles (n) × Molar Mass (M)
  • Number of Particles (N) = Number of moles (n) × Avogadro's Number (N<0xE2><0x82><0x90>)

• Exam Relevance: Extremely important. Expect numerical problems involving conversions between mass, moles, and number of particles. Understand the definitions clearly.

Example Calculation:
Calculate the number of moles in 88g of CO₂. (Atomic mass C=12u, O=16u)

1. Calculate Molar Mass (M) of CO₂ = 12 + (2 × 16) = 12 + 32 = 44 g/mol.
2. Given Mass (m) = 88g.
3. Number of moles (n) = m / M = 88g / 44 g/mol = 2 moles.

Calculate the number of molecules in 2 moles of CO₂.
Number of molecules (N) = n × N<0xE2><0x82><0x90> = 2 × 6.022 × 10²³ = 12.044 × 10²³ molecules.

Multiple Choice Questions (MCQs)

Here are 10 MCQs based on the chapter for your practice:

1. Which law states that mass can neither be created nor destroyed in a chemical reaction?
   a) Law of Definite Proportions
   b) Law of Multiple Proportions
   c) Law of Conservation of Mass
   d) Avogadro's Law

2. The ratio by mass of Hydrogen to Oxygen in water (H₂O) is always:
   a) 1:16
   b) 1:8
   c) 2:1
   d) 8:1

3. Which postulate of Dalton's atomic theory has been invalidated by the discovery of isotopes?
   a) Atoms are indivisible.
   b) Atoms of different elements have different masses.
   c) Atoms combine in small whole number ratios.
   d) Atoms of a given element are identical in mass and properties.

4. What is the chemical symbol for Iron?
   a) Ir
   b) I
   c) Fe
   d) Rn

5. The formula for Aluminium Sulphate is:
   a) AlSO₄
   b) Al₂(SO₄)₃
   c) Al₃(SO₄)₂
   d) Al(SO₄)₂

6. What is the atomicity of Ozone (O₃)?
   a) 1
   b) 2
   c) 3
   d) Polyatomic (but specifically 3)

7. Calculate the molecular mass of Sulphuric Acid (H₂SO₄). (Atomic masses: H=1u, S=32u, O=16u)
   a) 49u
   b) 98u
   c) 50u
   d) 100u

8. How many moles are present in 9 grams of water (H₂O)? (Molar mass of H₂O = 18 g/mol)
   a) 1 mole
   b) 2 moles
   c) 0.5 moles
   d) 9 moles

9. Avogadro's constant represents the number of:
   a) Grams in one mole
   b) Atoms in 12g of Carbon-12
   c) Moles in one gram
   d) Molecules in 22.4 L of a gas

10. Which of the following is a polyatomic cation?
    a) Sulphate (SO₄²⁻)
    b) Chloride (Cl⁻)
    c) Ammonium (NH₄⁺)
    d) Oxide (O²⁻)

Answers to MCQs:

1. c) Law of Conservation of Mass
2. b) 1:8
3. d) Atoms of a given element are identical in mass and properties.
4. c) Fe
5. b) Al₂(SO₄)₃
6. c) 3
7. b) 98u (Calculation: (2x1) + 32 + (4x16) = 2 + 32 + 64 = 98u)
8. c) 0.5 moles (Calculation: n = m/M = 9g / 18 g/mol = 0.5 mol)
9. b) Atoms in 12g of Carbon-12 (This is the basis of the definition of a mole and Avogadro's number)
10. c) Ammonium (NH₄⁺)

Study these notes thoroughly. Focus on understanding the concepts, especially the mole concept and formula writing, as they are frequently tested. Good luck with your preparation!