Atoms and Molecules

Laws of chemical combination

Law of conservation of mass

In a chemical reaction, mass can neither be created nor destroyed. Total mass of reactants = total mass of products. Antoine Lavoisier showed this. If a gas can escape, the balance seems to drop, so the test is done in a sealed flask.

Example: 5.3 g of sodium carbonate reacts with 6.0 g of ethanoic acid to give 2.2 g of carbon dioxide, 0.9 g of water and 8.2 g of sodium ethanoate. Reactants: 5.3 + 6.0 = 11.3 g. Products: 2.2 + 0.9 + 8.2 = 11.3 g. Equal.

Law of constant proportions (definite proportions)

In a pure compound, the elements are always present in the same proportion by mass, whatever its source or the way it is made. Joseph Proust stated it.

Water: hydrogen : oxygen = 1 : 8 by mass. 9 g of water has 1 g H and 8 g O; 36 g has 4 g H and 32 g O. Ammonia: nitrogen : hydrogen = 14 : 3. Carbon dioxide: carbon : oxygen = 3 : 8.

Dalton's atomic theory

John Dalton (1808) explained both laws with these points:

1. All matter is made of very tiny particles called atoms.
2. Atoms cannot be created or destroyed in a chemical reaction. (This explains conservation of mass.)
3. Atoms of one element are the same in mass and properties; atoms of different elements differ.
4. Atoms combine in small whole-number ratios to form compounds. (This explains constant proportions.)
5. In a given compound, the kinds and numbers of atoms are fixed.
6. Atoms were thought to be indivisible.

Limits we know today: atoms can be divided into electrons, protons and neutrons; atoms of the same element can have different masses (isotopes); atoms of different elements can have the same mass (isobars).

Atoms, symbols and atomic mass

An atom is the smallest particle of an element that takes part in a reaction. Atoms are about 10⁻¹⁰ m across. Each element has a symbol: H, O, C, N, Na (from Latin natrium), Fe (ferrum), K (kalium), Cu (cuprum). The atomic mass unit (u) is 1/12 of the mass of one carbon-12 atom. So H ≈ 1 u, C = 12 u, O = 16 u.

Molecules, ions, ionic and covalent compounds

A molecule is a group of two or more atoms joined together; it is the smallest particle of a substance that can exist on its own and shows its properties.

• Molecule of an element: same atoms. O₂, N₂, H₂, Cl₂ (atomicity 2), O₃ ozone (3), P₄ (4), S₈ (8). Noble gases like He and Ar exist as single atoms (atomicity 1).
• Molecule of a compound: different atoms in a fixed ratio. H₂O, CO₂, NH₃, CH₄.

Atomicity = number of atoms in one molecule.

An ion is an atom or a group of atoms with a charge. A positive ion is a cation (Na⁺, Mg²⁺, NH₄⁺); a negative ion is an anion (Cl⁻, O²⁻, SO₄²⁻). A group of atoms with a charge is a polyatomic ion.

Ionic compounds (usually a metal + a non-metal) are made of ions held by attraction: NaCl, MgO, CaCl₂. They do not have molecules; we speak of a formula unit. Covalent (molecular) compounds (non-metals only) are made of molecules in which atoms share electrons: H₂O, CO₂, CH₄.

Writing chemical formulae

Valency is the combining power of an atom or ion: how many electrons it gives, takes or shares. H = 1, Na = 1, Cl = 1, O = 2, Mg = 2, Ca = 2, Al = 3, N = 3 (in NH₃), C = 4.

Rules:

1. Write the symbols, metal / positive ion first.
2. Write the valency under each symbol.
3. Cross over the valencies; they become subscripts.
4. Cut down to the simplest ratio; do not write 1.
5. For a polyatomic ion used more than once, put it in brackets: Ca(OH)₂, (NH₄)₂SO₄.

Examples: H (1) + Cl (1) → HCl. Mg (2) + Cl (1) → MgCl₂. Al (3) + O (2) → Al₂O₃. Ca (2) + O (2) → Ca₂O₂ → CaO. Al (3) + SO₄ (2) → Al₂(SO₄)₃. The total positive charge always equals the total negative charge.

Molecular mass and formula unit mass

Molecular mass = sum of the atomic masses of all atoms in one molecule, in u.

H₂O = 2 × 1 + 16 = 18 u. CO₂ = 12 + 2 × 16 = 44 u. H₂SO₄ = 2 × 1 + 32 + 4 × 16 = 98 u.

For ionic compounds we use formula unit mass, found the same way from the formula: NaCl = 23 + 35.5 = 58.5 u. CaCO₃ = 40 + 12 + 3 × 16 = 100 u.

Method: list each element, multiply its atomic mass by how many atoms there are, then add. Use the free-play step in the 3D to watch it line by line.

Next step (higher classes): 1 mole of any substance has 6.022 × 10²³ particles, and its mass in grams equals its molecular mass (18 g of water is 1 mole).

Try it

1. Balance test at home (with an adult): put a spoon of baking soda in a small bottle and some vinegar in a balloon. Fit the balloon on the bottle and weigh it on a kitchen scale. Tip the vinegar in. The balloon fills with gas, but the reading does not change.
2. Formula game: write valency cards (Na 1, Mg 2, Al 3, Cl 1, O 2, SO₄ 2). Pick one positive and one negative card and cross them. Check your answer by making the charges add to zero.
3. Predict then check: before choosing H₂SO₄ in the 3D, work out its molecular mass on paper. Then compare with the line-by-line sum.

Key formulas and definitions

• Mass of reactants = mass of products
• Water: H : O = 1 : 8 by mass
• Molecular mass = Σ (number of atoms × atomic mass)
• Atomic masses (u): H 1, C 12, N 14, O 16, Na 23, Mg 24, Al 27, S 32, Cl 35.5, K 39, Ca 40
• Formula by criss-cross: A (valency a) + B (valency b) → A_b B_a, then simplify

Worked examples

Common mistakes

• Forgetting to simplify: Ca²⁺ + O²⁻ gives CaO, not Ca₂O₂.
• Missing brackets around a repeated polyatomic ion: write Ca(OH)₂, not CaOH₂.
• Multiplying only the first atom: in H₂SO₄, O is 4 × 16 = 64 u, not 16 u.
• Saying NaCl has molecules. It is ionic; use 'formula unit mass'.