GCSE Chemistry Equations & mass revision

A chemical equation is a sentence: reactants on the left, products on the right, an arrow meaning “reacts to form”. Writing one well demands correct formulae first. The traps to memorise: hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine exist as diatomic molecules (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂), and ionic compounds must be charge-neutral — calcium hydroxide is Ca(OH)₂ because two 1− hydroxides balance one 2+ calcium. Sodium plus water is “sodium + water → sodium hydroxide + hydrogen” as words; as formulae it becomes 2Na + 2H₂O → 2NaOH + H₂ once balanced.

Balancing adds coefficients — the large numbers in front of formulae — until each element has equal counts on both sides. A reliable method: balance metals first, then non-metals, then hydrogen, then oxygen (hydrogen and oxygen usually appear in several places, so save them for last). Check by listing every element’s totals. For CH₄ + O₂ → CO₂ + H₂O: carbon balances 1:1; four hydrogens need 2H₂O; those bring two oxygens, which with CO₂’s two make four, so 2O₂: CH₄ + 2O₂ → CO₂ + 2H₂O.

The coefficients also carry the ratio of moles: in 2Mg + O₂ → 2MgO, two moles of magnesium react with one of oxygen to give two of magnesium oxide. That ratio is the engine of all the reacting-mass work in the next chapter.

The law of conservation of mass says the total mass of products equals the total mass of reactants: reactions only rearrange atoms. In a sealed flask, a balance confirms it every time. In an open vessel, the balance can lie. When calcium carbonate fizzes with acid, CO₂ bubbles away — the flask loses exactly the mass of gas that escaped. When magnesium burns in air, oxygen joins the solid, and the ash of magnesium oxide outweighs the original ribbon. Neither reaction breaks the law; the system just traded matter with the air.

The reverse case matters too: burning 2.4 g of magnesium in a crucible can leave more than 2.4 g of white powder, because oxygen atoms from the air became part of the product. 2.4 g Mg contains 0.1 mol of Mg atoms; 0.1 mol of MgO has mass 4.0 g, and the extra 1.6 g is exactly the oxygen gained (0.1 mol × 16 g/mol).

The relative atomic mass (Ar) of an element is its average atomic mass compared with carbon-12 — the number printed on the periodic table. The relative formula mass (Mr) of a compound is the Ar values of its formula added together. Water: (2 × 1) + 16 = 18. Carbon dioxide: 12 + (2 × 16) = 44. Magnesium oxide: 24 + 16 = 40. Brackets multiply: Ca(OH)₂ = 40 + 2 × (16 + 1) = 74, and (NH₄)₂SO₄ = 2 × (14 + 4) + 32 + 4 × 16 = 132.

The empirical formula is the simplest whole-number ratio of atoms in a compound. From composition data the recipe is always the same three moves: divide each element’s mass (or percentage) by its Ar to get moles of atoms; divide every result by the smallest; if any ratio is not a whole number, multiply through to clear fractions (1.5 needs ×2, 1.33 ×3, 1.25 ×4). For a compound 40% calcium, 12% carbon and 48% oxygen: 40 ÷ 40 = 1, 12 ÷ 12 = 1, 48 ÷ 16 = 3 — ratio 1:1:3, empirical formula CaCO₃.

When solutions react, many dissolved ions take no part in the change — they are spectator ions. The neutralisation of hydroxide by acid in every sodium-hydroxide-plus-hydrochloric-acid beaker is really just H⁺(aq) + OH⁻(aq) → H₂O(l): the Na⁺ and Cl⁻ ride along unchanged. Writing an ionic equation: write the full balanced equation with state symbols, split only aqueous ionic compounds into their ions, then cancel the spectators. For precipitation — AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq) — the ionic equation is Ag⁺(aq) + Cl⁻(aq) → AgCl(s), a white precipitate forming out of solution.

A thermometer marked in 1 °C steps can justify a reading “to the nearest 1 °C”, but the true value may sit anywhere within ±0.5 °C: the uncertainty of a single reading is half the resolution of the instrument. Repeated measurements attack the problem differently: repeat a titration three times, discard the anomaly, and report the mean with the range — “24.35 cm³ (range 0.20 cm³)” — so the spread shows how much to trust the number. The tighter the range, the smaller the uncertainty of the mean.