Science and Technology Class 9 Notes Chapter 4 Maharashtra Board

Measurement of Matter

Introduction to Matter and Dalton’s Atomic Theory

What is Dalton’s Atomic Theory?

• Proposed by John Dalton in the early 1800s.

Main points:

1. Matter is made of tiny particles called atoms.
2. Atoms of the same element are identical; atoms of different elements are different.
3. Atoms combine in whole-number ratios to form compounds.
4. Atoms cannot be created or destroyed in a chemical reaction.

How Are Compounds Formed?

• Compounds form when atoms of different elements chemically combine in a fixed ratio.
• Example: Water (H₂O) forms when 2 hydrogen atoms combine with 1 oxygen atom.

Molecular Formulae of Common Compounds:

• Common Salt: NaCl (Sodium Chloride).
• Slaked Lime: Ca(OH)₂ (Calcium Hydroxide).
• Water: H₂O.
• Lime: CaO (Calcium Oxide).
• Limestone: CaCO₃ (Calcium Carbonate).

Laws of Chemical Combination

What Are They?

• These laws explain how elements combine to form compounds, discovered in the 18th and 19th centuries through experiments.

1. Law of Conservation of Matter:

• Proposed by Antoine Lavoisier in 1785.
• States: The total mass of reactants equals the total mass of products in a chemical reaction.

Example (Activity 1):

• • Mix 56 g of CaO (calcium oxide) with 18 g of H₂O (water).
  • Product: Ca(OH)₂ (slaked lime) = 56 + 18 = 74 g.
  • Mass before and after the reaction is the same (74 g), proving the law.

Example (Activity 2):

• • Mix solutions of CaCl₂ (calcium chloride) and Na₂SO₄ (sodium sulphate).
  • Reaction: CaCl₂ + Na₂SO₄ → CaSO₄ (precipitate) + 2NaCl.
  • Mass of the flask before and after mixing remains the same, confirming the law.

2. Law of Constant Proportion:

• Proposed by J.L. Proust in 1794.
• States: In a compound, the elements are always present in a fixed proportion by mass, no matter the source.

Example:

• • In water (H₂O), hydrogen and oxygen are always in the ratio 1:8 by mass (1 g H + 8 g O = 9 g water).
  • In CO₂, carbon and oxygen are in the ratio 3:8 (12 g C + 32 g O = 44 g CO₂).

Verification Experiment:

Two samples of copper oxide (CuO) were made:

• • One from CuCO₃ (copper carbonate).
  • Another from Cu(NO₃)₂ (copper nitrate).

Both 8 g samples of CuO were treated with hydrogen:

• • Result: 6.4 g copper (Cu) + 1.8 g water (H₂O).
  • Water (H₂O) has H:O = 1:8, so 1.8 g water has 1.6 g oxygen.
  • So, CuO has 6.4 g Cu and 1.6 g O → ratio = 4:1.

Expected ratio from CuO formula (Cu = 63.5, O = 16): 63.5:16 ≈ 4:1.

Experimental ratio matches the expected ratio, proving the law.

Atom: Size, Mass, and Valency

Structure of an Atom:

• Nucleus (center): Contains protons (positive) and neutrons (neutral).
• Extra-nuclear part: Electrons (negative) move around the nucleus in orbits.

Size of an Atom:

• Measured by atomic radius (distance from nucleus to outermost orbit).
• Unit: Nanometer (nm). 1 m = 10⁹ nm.
• Size depends on:
  • More orbits → larger atom (e.g., K is larger than Na).
  • Same outermost orbit but more electrons → smaller atom (e.g., Mg is smaller than Na).
• Seen using modern tools like field ion microscopes (e.g., iridium atoms in Figure 4.2).

Mass of an Atom:

• Mass is mostly in the nucleus (due to protons and neutrons).
• Atomic Mass Number = Protons + Neutrons (called nucleons).

Relative Atomic Mass:

• • Initially, hydrogen (H) was the reference (mass = 1, as it has 1 proton).
  • Example: Nitrogen (N) is 14 times heavier than H, so its relative mass = 14.
  • Later, carbon (C) was chosen as the reference in 1961 (mass = 12).
  • On this scale, H = 1 (12 × 1/12).

Unified Atomic Mass (Dalton, u):

• • Now used instead of relative mass.
  • 1 u = 1.66053904 × 10⁻²⁷ kg.
  • Example: Carbon = 12 u, Hydrogen = 1 u.

Valency:

• Valency is the combining capacity of an element (how many bonds it can form).
• It’s the number of electrons an atom gives or takes to form an ionic bond.

Example (NaCl):

• • Sodium (Na): 2,8,1 → gives 1 electron → Na⁺ (2,8) → valency = 1.
  • Chlorine (Cl): 2,8,7 → takes 1 electron → Cl⁻ (2,8,8) → valency = 1.
  • Na⁺ and Cl⁻ form NaCl through ionic bonding.

Example (MgCl₂):

• • Mg: 2,8,2 → gives 2 electrons → Mg²⁺ → valency = 2.
  • Cl: Takes 1 electron each → 2 Cl⁻ → MgCl₂.

Example (CaO):

• • Ca: 2,8,8,2 → gives 2 electrons → Ca²⁺ → valency = 2.
  • O: 2,6 → takes 2 electrons → O²⁻ → valency = 2 → CaO.

Variable Valency:

• Some elements show different valencies under different conditions.
• Example: Iron (Fe) has valency 2 in FeCl₂ and 3 in FeCl₃.

Molecules of Elements and Compound

Molecules of Elements:

• Monoatomic: Exist as single atoms (e.g., He, Ne).
• Diatomic: Two atoms combine (e.g., O₂, N₂).

Molecules of Compounds:

• Formed when atoms of different elements combine.

Example: H₂O (2 H + 1 O), NaCl (1 Na + 1 Cl).

Molecular Mass and the Concept of Mole

Molecular Mass:

• The sum of atomic masses of all atoms in a molecule.
• Unit: Dalton (u).
• Example (H₂O):
  • H = 1 u, O = 16 u.
  • H₂O = (2 × 1) + 16 = 18 u.

Examples:

• • NaCl: 23 + 35.5 = 58.5 u.
  • MgCl₂: 24 + (2 × 35.5) = 95 u.
  • KNO₃: 39 + 14 + (3 × 16) = 101 u.
  • H₂SO₄: (2 × 1) + 32 + (4 × 16) = 98 u.

Mole Concept:

• A mole is the amount of a substance whose mass in grams equals its molecular mass in Daltons.

Example:

• • Molecular mass of O₂ = 32 u → 1 mole of O₂ = 32 g.
  • Molecular mass of H₂O = 18 u → 1 mole of H₂O = 18 g.

Formula: Number of moles (n) = Mass (g) / Molecular Mass.

Avogadro’s Number (Nₐ):

• • 1 mole of any substance has 6.022 × 10²³ molecules.
  • Example: 18 g of H₂O (1 mole) has 6.022 × 10²³ molecules.

Example: How many molecules in 66 g of CO₂?

• • Molecular mass of CO₂ = 12 + (2 × 16) = 44 u.
  • Moles = 66 / 44 = 1.5 moles.
  • Molecules = 1.5 × 6.022 × 10²³ = 9.033 × 10²³.

Radicals

What Are Radicals?

• Ions (charged particles) in ionic compounds that act independently in reactions.

Two types:

• • Cations (positive, basic radicals): e.g., Na⁺, Ca²⁺, NH₄⁺.
  • Anions (negative, acidic radicals): e.g., Cl⁻, SO₄²⁻, NO₃⁻.

Basic Radicals: Form bases (e.g., NaOH, KOH).

Acidic Radicals: Form acids (e.g., HCl, H₂SO₄).

Simple Radicals: Single atoms (e.g., Na⁺, Cl⁻).

Composite Radicals: Group of atoms (e.g., SO₄²⁻, NH₄⁺).

Examples:

• NaOH: Na⁺ (basic) + OH⁻ (acidic).
• H₂SO₄: 2H⁺ (basic) + SO₄²⁻ (acidic).

Classify:

• • Basic: Ag⁺, Mg²⁺, Fe²⁺, NH₄⁺, Ca²⁺, K⁺, Na⁺.
  • Acidic: Cl⁻, SO₄²⁻, ClO₃⁻, NO₃⁻, Br⁻.

Writing Chemical Formulae

Steps (Cross-Multiplication Method):

1. Write the symbols of radicals (basic on the left, acidic on the right).
2. Write their valencies below.
3. Cross-multiply the valencies to get the number of each ion.
4. Write the final formula (omit charges, use subscripts).

• • Example (Na₂SO₄):
    • Na⁺ (valency 1), SO₄²⁻ (valency 2).
    • Cross-multiply: Na₂SO₄ (2 Na⁺ for 1 SO₄²⁻).

Examples:

• • Calcium Carbonate: Ca²⁺ + CO₃²⁻ → CaCO₃.
  • Sodium Bicarbonate: Na⁺ + HCO₃⁻ → NaHCO₃.
  • Silver Chloride: Ag⁺ + Cl⁻ → AgCl.
  • Magnesium Oxide: Mg²⁺ + O²⁻ → MgO.