Atomic Foundations of Matter Class 9 Notes and Questions

These Atomic Foundations of Matter Class 9 notes follow the chapter in order: the two laws of chemical combination, Dalton's atomic theory, covalent and ionic bonding, naming compounds, writing formulae, and calculating molecular and formula unit mass. Every calculation is worked out step by step, and the important questions at the end have model answers.

Law of conservation of mass

Proposed by Antoine Lavoisier in 1789: matter can neither be created nor destroyed in a chemical reaction. The total mass of the reactants equals the total mass of the products.

• In a physical change, like dissolving salt in water, the solution's mass equals water + salt.
• In a reaction that gives off a gas, the mass may seem to drop if the gas escapes. In a closed system, the mass stays the same.
• Mixing sodium sulfate and barium chloride solutions gives a white precipitate of barium sulfate, plus sodium chloride. The total mass before and after is identical, because atoms only rearrange.

Worked example (from the chapter): calcium carbonate (4.0 g) + hydrochloric acid (2.92 g) → calcium chloride (4.44 g) + water (0.72 g) + carbon dioxide (1.76 g).

• Reactants: 4.0 + 2.92 = 6.92 g
• Products: 4.44 + 0.72 + 1.76 = 6.92 g
• Equal, so the law is obeyed.

Law of constant proportions

Proposed by Joseph Proust: a pure compound always contains the same elements in a fixed ratio by mass, whatever its source or method of preparation.

• Water from a river, a well or a lab always has hydrogen and oxygen in the mass ratio 1 : 8. So 9 g of water always gives 1 g of hydrogen and 8 g of oxygen.
• Proust found copper carbonate always had copper, carbon and oxygen in the same proportions, natural or lab-made.
• If you take elements in a different ratio, the extra amount of one element stays unreacted. This is how compounds differ from mixtures, which can have any ratio.

Worked example: in sodium chloride, sodium : chlorine = 23 : 35.5 by mass. 46 g of sodium is 23 × 2, so it needs 35.5 × 2 = 71 g of chlorine.

Dalton's atomic theory

John Dalton gave his theory in 1808 to explain these laws. A postulate is a basic assumption accepted without formal proof. Key postulates:

• All matter is made of tiny, indivisible particles called atoms.
• Atoms cannot be created or destroyed in a chemical reaction. This explains conservation of mass.
• Atoms of one element are identical in mass and properties; atoms of different elements differ.
• Atoms combine in the ratio of simple whole numbers, and in a given compound the relative number and kinds of atoms are constant. This explains constant proportions.
• Reactions only separate, join or rearrange atoms.

Dalton's ideas helped turn chemistry from a qualitative study into a quantitative science.

Covalent bonding and naming

A molecule is an electrically neutral group of two or more atoms held by chemical bonds. Atoms with fewer than eight valence electrons (or two, for the K-shell) are unstable, so they share electrons to reach an octet or duplet.

• One shared pair: single bond (H—H).
• Two shared pairs (four electrons): double bond (O=O).
• Covalent bonds usually form between non-metals. In H₂O, oxygen shares electrons with two hydrogen atoms; oxygen completes its octet and each hydrogen its duplet.

Naming binary covalent compounds: keep the first element's name; change the second to end in "-ide"; use prefixes mono (1), di (2), tri (3), tetra (4), penta (5), hexa (6).

• "Mono" is usually left out for the first element.
• If a prefix ends in "a" or "o" and the element starts with a vowel, drop that letter: monoxide, tetroxide. Prefixes ending in "i" keep it.
• When hydrogen comes first, prefixes are generally not used: H₂S is hydrogen sulfide. Some have common names: water (H₂O), ammonia (NH₃).

Ionic bonding

Metals with few valence electrons lose them to form positive cations. Non-metals with many valence electrons gain electrons to form negative anions. The strong electrostatic attraction between them is the ionic bond.

• Sodium (2, 8, 1) loses one electron → Na⁺.
• Chlorine (2, 8, 7) gains that electron → Cl⁻.
• Together they form neutral NaCl.

Ionic compounds do not exist as separate molecules. They form 3-D repeating patterns called crystal lattices. In NaCl, each sodium ion is surrounded by six chloride ions, and each chloride ion by six sodium ions.

Writing formulae: the criss-cross method

1. Write the symbols, cation (metal) first.
2. Write each one's valency or charge below it.
3. Criss-cross the numbers to become the other symbol's subscript.
4. Simplify to the lowest whole-number ratio; never write a subscript of 1.
5. Put a polyatomic ion (like OH⁻ or SO₄²⁻) in brackets if you need more than one.

Examples:

• Mg²⁺ and O²⁻: Mg₂O₂ simplifies to MgO.
• Ca²⁺ and Cl⁻: Ca₁Cl₂ → CaCl₂.
• Mg²⁺ and OH⁻: Mg(OH)₂.
• Al³⁺ and SO₄²⁻: the 2 goes to Al and the 3 to sulfate → Al₂(SO₄)₃.

Properties, and molecular mass

Molecular mass (covalent) and formula unit mass (ionic) are both found the same way: multiply each atom's atomic mass (in u) by its number in the formula, then add.

Worked example: Ca(NO₃)₂ with Ca = 40 u, N = 14 u, O = 16 u.

• One NO₃ = 14 + (16 × 3) = 14 + 48 = 62 u
• Two NO₃ = 62 × 2 = 124 u
• Total = 40 + 124 = 164 u

Remember this

• Mass of reactants = mass of products.
• Water: hydrogen : oxygen = 1 : 8 by mass.
• Covalent = sharing; ionic = transfer.
• Criss-cross, then simplify; brackets for polyatomic ions.
• Ionic solids do not conduct; their solutions and melts do.

Important questions with answers

1. 12 g of carbon reacts with 32 g of oxygen. What mass of carbon dioxide forms? 12 + 32 = 44 g, by the law of conservation of mass.

2. How much hydrogen and oxygen are in 18 g of water? The ratio is 1 : 8, which is 9 parts. 18 ÷ 9 = 2, so hydrogen = 2 g and oxygen = 16 g.

3. Student A makes copper oxide with copper : oxygen = 4 : 1, and student B with 8 : 2. Does this follow the law of constant proportions? Yes. 8 : 2 simplifies to 4 : 1, the same ratio.

4. Which of Dalton's postulates explains the law of constant proportions? Atoms combine in simple whole-number ratios, and the relative number and kinds of atoms in a compound are constant.

5. How many electrons are shared in a double bond? Two pairs, so four electrons.

6. Write the formula of aluminium oxide (Al³⁺, O²⁻). Criss-cross gives Al₂O₃.

7. Calculate the molecular mass of HNO₃ (H = 1 u, N = 14 u, O = 16 u). 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 u.

8. Calculate the molecular mass of NH₃ (N = 14 u, H = 1 u). 14 + (3 × 1) = 17 u.

9. Why does salt solution conduct electricity but sugar solution does not? NaCl splits into free Na⁺ and Cl⁻ ions in water that carry current. Sugar is covalent and stays as neutral molecules.

Common mistakes to avoid

• Writing Mg₂O₂ instead of simplifying to MgO.
• Forgetting brackets: write Mg(OH)₂, not MgOH₂.
• Calling NaCl a molecule; it forms a crystal lattice.
• Saying solid salt conducts electricity.

When you want to practise criss-cross formulae and mass sums with feedback, study this chapter with Joy.