Class 9 Science Chapter 9: Atomic Foundations of Matter — Notes, Examples & Practice

Law of conservation of mass (Lavoisier): in a chemical reaction, mass is neither created nor destroyed — the total mass of the products equals the total mass of the reactants (in a closed system, so no gas escapes).

Law of constant (definite) proportions (Proust): a pure chemical compound always contains the same elements combined in the same proportion by mass. Water is always hydrogen : oxygen = 1 : 8 by mass; carbon dioxide is always C : O = 3 : 8.

Dalton's atomic theory (1808) explained these laws: matter is made of tiny particles called atoms; atoms of an element are identical in mass and properties; atoms combine in small whole-number ratios to form compounds; atoms are neither created nor destroyed in reactions, only rearranged. (We now know atoms are divisible, and isotopes have different masses.)

An atom is the smallest particle of an element that takes part in a chemical reaction. Atoms are extremely small — a hydrogen atom has a radius of about $10^{-10}$ m. Each element has a symbol of one or two letters: H, O, C, Na (from Latin natrium), Fe (ferrum), K (kalium).

A molecule is the smallest particle of an element or compound that can exist independently and shows its properties. It is a group of two or more atoms chemically bonded: O₂, H₂O, CO₂.

Atomicity is the number of atoms in one molecule of an element: monatomic (He, Ne, Ar — noble gases), diatomic (H₂, O₂, N₂, Cl₂), triatomic (O₃, ozone), tetra-atomic (P₄), polyatomic (S₈). Metals and carbon are written simply as Fe, Cu, C because they form large structures.

An ion is a charged particle formed when an atom (or group of atoms) loses or gains electrons. Losing electrons gives a positive ion, a cation (Na⁺, Mg²⁺, mostly metals). Gaining electrons gives a negative ion, an anion (Cl⁻, O²⁻, mostly non-metals). Groups of atoms carrying a charge are polyatomic ions: NH₄⁺ (ammonium), OH⁻ (hydroxide), NO₃⁻ (nitrate), SO₄²⁻ (sulphate), CO₃²⁻ (carbonate).

Ionic bonding: a metal atom transfers electrons to a non-metal atom; the oppositely charged ions attract strongly. Sodium (2, 8, 1) gives one electron to chlorine (2, 8, 7), making Na⁺ (2, 8) and Cl⁻ (2, 8, 8). Ionic compounds are usually hard solids with high melting points, and they conduct electricity when molten or dissolved in water (ions are free to move), but not as solids.

Covalent bonding: non-metal atoms share pairs of electrons to complete their outer shells: H₂, Cl₂, H₂O, CH₄. Covalent compounds are made of molecules, often have low melting and boiling points and usually do not conduct electricity.

Valency is the combining capacity of an atom or ion, equal to the size of its charge for ions: Na⁺ 1, Ca²⁺ 2, Al³⁺ 3, Cl⁻ 1, O²⁻ 2, N³⁻ 3; NH₄⁺ 1, OH⁻ 1, NO₃⁻ 1, SO₄²⁻ 2, CO₃²⁻ 2, PO₄³⁻ 3.

A compound is electrically neutral: total positive charge = total negative charge. Criss-cross method: write the symbols (cation first), write their valencies below, and swap them as subscripts. Then reduce to the simplest whole-number ratio, and omit the subscript 1.

Use brackets around a polyatomic ion when there are two or more of it: Ca(OH)₂, Al₂(SO₄)₃. A subscript (H₂O) tells the number of atoms in the formula; a coefficient in front (2H₂O) tells the number of molecules — they are not interchangeable.

Atoms are too light to weigh in grams conveniently, so we use relative atomic mass on a scale where one atom of carbon-12 has a mass of exactly 12 atomic mass units (u). So 1 u is $\dfrac{1}{12}$ of the mass of a C-12 atom. Common values: H 1, C 12, N 14, O 16, Na 23, Mg 24, S 32, Cl 35.5, Ca 40.

Molecular mass is the sum of the atomic masses of all atoms in a molecule: H₂O $= 2(1) + 16 = 18$ u.

For ionic compounds, which are not made of molecules, we calculate the formula unit mass the same way: NaCl $= 23 + 35.5 = 58.5$ u.