Ans: The Law of Conservation of Mass states that in a chemical reaction, the total mass of reactants is equal to the total mass of products. Mass is neither created nor destroyed during a chemical reaction. (Given by Antoine Lavoisier, 1789.)

Ans: An atom is the smallest particle of an element that can take part in a chemical reaction. It may or may not exist independently. The SI unit of atomic mass is unified atomic mass unit (u), where 1u = 1/12 the mass of one atom of carbon-12 = 1.66 × 10⁻²⁴ g.

Ans: Law of Constant Proportions (Proust's Law): In a pure chemical compound, elements are always present in definite proportions by mass. Example: Water (H₂O) always contains H and O in mass ratio 1:8 regardless of its source. CO₂ always has C:O in ratio 3:8. This law shows that compounds have a fixed, definite composition.

Ans: A mole is defined as the amount of a substance that contains as many elementary entities (atoms, molecules, ions) as there are atoms in exactly 12 g of carbon-12. This number is called Avogadro's number (Nₐ) = 6.022 × 10²³. So 1 mole of any substance contains 6.022 × 10²³ entities. Molar mass = atomic/molecular mass expressed in grams = mass of 1 mole of the substance.

Ans: Atom: Smallest particle of an element; may not exist independently (e.g., Na, O); has atomic mass. Molecule: Smallest particle of a substance (element or compound) that can exist independently; formed by combination of two or more atoms; has molecular mass = sum of atomic masses. Examples: O₂, N₂, H₂O, CO₂. Diatomic molecules: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂. Monoatomic molecules: Noble gases (He, Ne, Ar).

Ans: (a) Water: H₂O. Mol. mass = 2(1) + 16 = 18 u. (b) Carbon dioxide: CO₂. Mol. mass = 12 + 2(16) = 44 u. (c) Glucose: C₆H₁₂O₆. Mol. mass = 6(12) + 12(1) + 6(16) = 72 + 12 + 96 = 180 u. (d) Sulphuric acid: H₂SO₄. Mol. mass = 2(1) + 32 + 4(16) = 2 + 32 + 64 = 98 u. Atomic masses used: H=1, C=12, N=14, O=16, S=32, Na=23, Cl=35.5.

Ans: Polyatomic ions are groups of atoms carrying a net electric charge, that act as a single unit in a compound. Examples: (1) Ammonium ion: NH₄⁺ (+1). (2) Hydroxide ion: OH⁻ (–1). (3) Carbonate ion: CO₃²⁻ (–2). (4) Sulphate ion: SO₄²⁻ (–2). (5) Nitrate ion: NO₃⁻ (–1). (6) Phosphate ion: PO₄³⁻ (–3). (7) Bicarbonate/Hydrogen carbonate: HCO₃⁻ (–1). These ions do not exist independently — they are found in ionic compounds such as Na₂CO₃, (NH₄)₂SO₄ etc.

Ans: (a) Molar mass of Na = 23 g/mol. Moles = 46/23 = 2 moles. (b) Molar mass of CO₂ = 44 g/mol. Mass = 2 × 44 = 88 g. (c) Molar mass of H₂O = 18 g/mol. Moles = 9/18 = 0.5 moles. Molecules = 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules. (d) Moles of O₂ = 0.5. Each O₂ has 2 atoms. Atoms = 0.5 × 2 × 6.022 × 10²³ = 6.022 × 10²³ atoms. Key formulae: Moles = Given mass/Molar mass; Number of particles = Moles × Nₐ; Molar mass = sum of atomic masses in grams.

Ans: Formula writing rule: Write the symbols of the ions with their valencies (charges). Cross-multiply the valencies (use criss-cross method) to get the subscripts. Simplify if possible. Calcium ion: Ca²⁺; Phosphate ion: PO₄³⁻. Cross-multiplying: Ca₃(PO₄)₂. Formula of calcium phosphate = Ca₃(PO₄)₂. Other examples: AlCl₃ (Al³⁺, Cl⁻); Na₂SO₄ (Na⁺, SO₄²⁻); MgO (Mg²⁺, O²⁻).

Ans: Atomicity is the number of atoms in one molecule of an element. Ozone (O₃): atomicity = 3. Phosphorus (P₄): atomicity = 4. Sulphur (S₈): atomicity = 8. Noble gases (He, Ne, Ar): atomicity = 1 (monoatomic). O₂, N₂, H₂: atomicity = 2 (diatomic).

Ans: Molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass units (u). Example: H₂O = 2 × 1 + 16 = 18 u.

Ans: Formula unit mass is the sum of the atomic masses of all the atoms in a formula unit of an ionic compound, expressed in u. Example: NaCl = 23 + 35.5 = 58.5 u.

Ans: Molecular mass of HNO₃ = 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 u.

Ans: Moles = 52 ÷ 4 = 13 mol. Helium is monoatomic, so number of atoms = 13 × 6.022 × 10²³ = 7.83 × 10²⁴ atoms.

Ans: An ion is a charged particle formed when an atom or group of atoms gains or loses electrons. A cation is a positively charged ion formed by loss of electrons, e.g. Na⁺. An anion is a negatively charged ion formed by gain of electrons, e.g. Cl⁻.

Ans: In water (H₂O), hydrogen and oxygen are always combined in the ratio 1 : 8 by mass (2 g of hydrogen to 16 g of oxygen). Water from any source has this same ratio, which illustrates the Law of Constant Proportions (Proust).

Ans: Moles = number of atoms ÷ Avogadro's number = 2.58 × 10²⁴ ÷ 6.022 × 10²³ = 4.28 mol. So the sample contains about 4.28 moles of oxygen atoms.

Ans: (a) Molar mass of N₂ = 28 g/mol, so mass = 0.5 × 28 = 14 g. (b) Molecules = 0.5 × 6.022 × 10²³ = 3.011 × 10²³. (c) Molar mass of H₂O = 18 g/mol, so moles = 9 ÷ 18 = 0.5 mol. (d) Molecules of H₂O = 0.5 × 6.022 × 10²³ = 3.011 × 10²³; each molecule has 3 atoms, so atoms = 3 × 3.011 × 10²³ = 9.033 × 10²³.

Ans: Mass of reactants = 5.3 + 6.0 = 11.3 g. Mass of products = 2.2 + 8.2 + 0.9 = 11.3 g. Since the total mass of reactants equals the total mass of products, mass is neither created nor destroyed in the reaction. This verifies the Law of Conservation of Mass (Lavoisier).

Ans: Postulates: (1) All matter is made of extremely small indivisible particles called atoms. (2) Atoms of a given element are identical in mass and properties. (3) Atoms of different elements have different masses and properties. (4) Atoms combine in small whole-number ratios to form compounds. (5) Atoms are neither created nor destroyed in a chemical reaction. Limitations: atoms are divisible into protons, neutrons and electrons; atoms of the same element can have different masses (isotopes); atoms of different elements can have the same mass (isobars).