Class 9 Science Chapter 3: Important Questions with Answers

Maharishi Kanad.

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

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

Manchester (1793).

About 10⁻¹⁰ m.

International Union of Pure and Applied Chemistry.

Au.

Carbon-12.

40 u.

2 (diatomic, Cl₂).

Na⁺ and Cl⁻.

HCO₃⁻.

A symbolic representation of the composition of a compound.

18 u.

The Italian scientist Amedeo Avogadro.

Both said that dividing matter ends in indivisible particles — Parmanu (Kanad), atoms (Democritus). Limitation: these were philosophical ideas, not tested by experiment till the eighteenth century.

Y in the flask, X in an ignition tube, cork and weigh; tilt to mix, a reaction occurs; weigh again — same mass. Conclusion: conservation of mass.

All matter is made of atoms; atoms are indivisible, neither created nor destroyed; atoms of an element are identical in mass and properties; atoms combine in small whole-number ratios.

First letter of the name (H, C); first two letters (Al, Co); first and a later letter (Cl, Zn); from Latin/German/Greek names (Fe, Na, K).

Earlier 1/16 of naturally occurring oxygen was the unit — oxygen forms compounds with many elements and gave mostly whole-number masses. In 1961 carbon-12 was chosen for a universally accepted unit: 1 u = 1/12 of C-12.

The seller calls a watermelon 12 units, cuts 12 pieces and weighs each fruit against one piece. Likewise each atom's mass is given relative to 1/12 of a C-12 atom.

An atom is the smallest particle of an element and mostly cannot exist independently; a molecule is a bonded group of two or more atoms that can exist independently and shows all properties of the substance.

Positively charged ion = cation (Na⁺, Ca²⁺, NH₄⁺); negatively charged = anion (Cl⁻, O²⁻, SO₄²⁻).

Valencies/charges must balance; metal first, non-metal after; more than one polyatomic ion goes in brackets.

Two hydroxide (OH) groups are joined to one calcium; atoms = 1 Ca + 2 O + 2 H = 5.

Both are sums of atomic masses; molecular mass for substances of molecules (H₂O 18 u), formula unit mass for substances of ions (NaCl 58.5 u).

2 molecules H₂ + 1 molecule O₂ → 2 molecules H₂O; 4 u H₂ + 32 u O₂ → 36 u H₂O.

(1) Conservation of mass: mass is neither created nor destroyed — Activity 3.1. (2) Constant proportions (Proust): elements in a compound are in a fixed mass ratio — water 1 : 8, ammonia 14 : 3. Dalton: atoms are neither created nor destroyed (→ law 1); the relative number and kinds of atoms are constant, in small whole numbers (→ law 2).

In CO, 3 g C + 4 g O; with C = 1, O = 1.33 — not whole. Earlier 1/16 of natural O was the unit (O forms many compounds, gives whole numbers). In 1961 carbon-12 became the standard: 1 u = 1/12 of one C-12 atom. Relative atomic mass = average atomic mass ÷ (1/12 of C-12). Fruit-seller analogy (Fig. 3.4).

Molecule: smallest particle that exists independently — O₂, H₂O. Atomicity: number of atoms in a molecule — He, Ar 1; O₂, N₂ 2; O₃ 3; P₄ 4; S₈ poly. Ion: a charged particle — Na⁺ (cation), Cl⁻ (anion); polyatomic ions — NH₄⁺, SO₄²⁻, PO₄³⁻. Metal–non-metal compounds are made of ions.

Al³⁺, O²⁻ → Al₂O₃; Ca²⁺, O²⁻ → Ca₂O₂ → CaO (simplify); Na⁺, NO₃⁻ → NaNO₃; Ca²⁺, OH⁻ → Ca(OH)₂ (bracket, two OH); NH₄⁺, SO₄²⁻ → (NH₄)₂SO₄.

1 mole = 6.022 × 10²³ particles (N₀) = relative mass in grams. n = m/M (52 g He → 13 mol); n = N/N₀ (12.044 × 10²³ He → 2 mol); m = M × n (0.5 mol N₂ → 14 g); N = (m/M) × N₀ (46 g Na → 12.044 × 10²³ atoms).

H₂ 2 u; O₂ 32 u; Cl₂ 71 u; CO₂ 12 + 32 = 44 u; CH₄ 12 + 4 = 16 u; NH₃ 14 + 3 = 17 u; ZnO 65 + 16 = 81 u; Na₂O 46 + 16 = 62 u.

Mass of reactants = 5.3 + 6 = 11.3 g; mass of products = 2.2 + 0.9 + 8.2 = 11.3 g. They are equal, so mass is conserved.

1 g H needs 8 g O, so 3 g H needs 3 × 8 = 24 g of oxygen.

Postulate (ii): atoms are indivisible particles which cannot be created or destroyed in a chemical reaction.

Postulate (vi): the relative number and kinds of atoms are constant in a given compound (with (v): atoms combine in small whole-number ratios).

A mass unit equal to exactly one-twelfth of the mass of one atom of carbon-12 = 1 u.

Atoms are extremely small (radius about 10⁻¹⁰ m, measured in nanometres) and atoms of most elements do not exist independently.

(i) Na₂O (ii) AlCl₃ (iii) Na₂S (iv) Mg(OH)₂

(i) aluminium sulphate (ii) calcium chloride (iii) potassium sulphate (iv) potassium nitrate (v) calcium carbonate

The symbolic representation of the composition of a compound — it shows the constituent elements and the number of atoms of each.

(i) 3 (2 H + 1 S) (ii) 5 (1 P + 4 O)

H₂ 2 u; O₂ 32 u; Cl₂ 71 u; CO₂ 44 u; CH₄ 16 u; C₂H₆ 30 u; C₂H₄ 28 u; NH₃ 17 u; CH₃OH 32 u.

ZnO = 65 + 16 = 81 u; Na₂O = 2 × 23 + 16 = 62 u; K₂CO₃ = 2 × 39 + 12 + 3 × 16 = 138 u.

12 g ÷ 6.022 × 10²³ = 1.99 × 10⁻²³ g.

Na: 100/23 = 4.35 mol; Fe: 100/56 = 1.79 mol. More moles = more atoms — 100 g of sodium has more atoms.

Boron = 0.096/0.24 × 100 = 40%; oxygen = 0.144/0.24 × 100 = 60%.

Still 11.00 g CO₂ — 3 g carbon combines with only 8 g oxygen; the remaining 42 g oxygen is left unused. Law of constant (definite) proportions.

Clusters of atoms carrying a fixed net charge that act as one ion — NH₄⁺, OH⁻, NO₃⁻, HCO₃⁻, CO₃²⁻, SO₄²⁻, PO₄³⁻.

(a) MgCl₂ (b) CaO (c) Cu(NO₃)₂ (d) AlCl₃ (e) CaCO₃

(a) CaO: calcium, oxygen (b) HBr: hydrogen, bromine (c) sodium hydrogen carbonate NaHCO₃: sodium, hydrogen, carbon, oxygen (d) K₂SO₄: potassium, sulphur, oxygen.

(a) 2 × 12 + 2 = 26 g (b) 8 × 32 = 256 g (c) 4 × 31 = 124 g (d) 1 + 35.5 = 36.5 g (e) 1 + 14 + 48 = 63 g (per mole).

(a) 14 g (b) 4 × 27 = 108 g (c) M = 46 + 32 + 48 = 126 g; 10 × 126 = 1260 g.

(a) 12/32 = 0.375 mol (b) 20/18 = 1.11 mol (c) 22/44 = 0.5 mol.

(a) 0.2 × 16 = 3.2 g (b) 0.5 × 18 = 9 g.

M(S₈) = 256 g; n = 16/256 = 0.0625 mol; N = 0.0625 × 6.022 × 10²³ = 3.76 × 10²² molecules.

M(Al₂O₃) = 54 + 48 = 102 g; n = 0.051/102 = 5 × 10⁻⁴ mol; each formula unit has 2 Al³⁺, so ions = 2 × 5 × 10⁻⁴ × 6.022 × 10²³ = 6.022 × 10²⁰.