Symbols, Formulae and Equations

Chemical symbols and formulae

Every element has a one- or two-letter symbol (first letter always capital). Compounds are represented by formulae showing the ratio of atoms. Key formulae to know:

• Diatomic elements (exist as pairs of atoms): H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.
• Common compounds: H₂O, CO₂, NH₃, CH₄, NaCl, CaCO₃, H₂SO₄, HCl, NaOH.

Writing ionic formulae

To write the formula of an ionic compound:

1. Write the cation then the anion.
2. Balance charges so the overall compound is neutral.
3. Simplify the ratio to the lowest whole numbers.

Example — aluminium oxide: Al³⁺ and O²⁻. LCM of 3 and 2 is 6 → need 2 Al³⁺ and 3 O²⁻ → Al₂O₃.

Common polyatomic ions to know:

• OH⁻ (hydroxide), NO₃⁻ (nitrate), SO₄²⁻ (sulfate), CO₃²⁻ (carbonate), NH₄⁺ (ammonium), PO₄³⁻ (phosphate).

State symbols

Add state symbols in equations: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).

Balancing equations

A balanced equation has the same number of atoms of each element on both sides. Only change coefficients (the big numbers in front of formulae) — never change subscripts.

Method:

1. Write the unbalanced equation with correct formulae.
2. Count atoms on each side.
3. Balance metals first, then non-metals, then hydrogen, then oxygen.
4. Check your work.

Example: iron + oxygen → iron(III) oxide Unbalanced: Fe + O₂ → Fe₂O₃ Balanced: 4Fe + 3O₂ → 2Fe₂O₃ Check: Fe: 4=4 ✓, O: 6=6 ✓.

Ionic equations

For reactions in solution, ionic equations show only the species that actually change. Spectator ions (present but unchanged) are omitted.

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) Full ionic: H⁺(aq) + Cl⁻(aq) + Na⁺(aq) + OH⁻(aq) → Na⁺(aq) + Cl⁻(aq) + H₂O(l) Net ionic: H⁺(aq) + OH⁻(aq) → H₂O(l)

The mole concept

One mole of any substance contains 6.02 × 10²³ particles (Avogadro's number, Nₐ). This is the link between the mass of a substance and the number of particles.

Molar mass M: the mass of one mole in grams = the relative formula mass in g/mol.

Key equations:

• n = m ÷ M (moles = mass ÷ molar mass)
• m = n × M
• M = m ÷ n

Example: How many moles are in 18 g of water (M = 18 g/mol)? n = 18 ÷ 18 = 1 mol.

Relative formula mass (Mᵣ): sum of relative atomic masses (Aᵣ) of all atoms in the formula. Example: Mᵣ of H₂SO₄ = 2(1) + 32 + 4(16) = 2 + 32 + 64 = 98.

Using moles in equations

The coefficients in a balanced equation give the molar ratio. Example: N₂ + 3H₂ → 2NH₃ — 1 mol N₂ reacts with 3 mol H₂ to give 2 mol NH₃.

If 0.5 mol N₂ reacts, it produces 1.0 mol NH₃.