This mini-lesson walks you through the whole of CCEA Quantitative Chemistry: relative formula mass, the mole, reacting masses, percentage composition and yield, empirical formulae, water of crystallisation, concentration, titration, gas volumes and atom economy.
Work through each screen, answer the questions as you go (some are concepts, most are calculations) and collect ⭐ stars. Use the periodic table values H=1, C=12, N=14, O=16, Na=23, Mg=24, S=32, Cl=35.5, Ca=40, Fe=56. Press Start when you're ready.
The relative formula mass (Mr) — called relative molecular mass for covalent molecules — is the sum of the relative atomic masses (Ar) of every atom in the formula.
Ca: 1 × 40 = 40
N: 2 × 14 = 28 (two NO₃ groups → two N)
O: 6 × 16 = 96 (two NO₃ groups → 2 × 3 = 6 O)
Mr = 40 + 28 + 96 = 164
Misconception: Mr has no units — it is a relative mass (a ratio to ¹²C), not a mass in grams. Only when you weigh out one mole does that number become a mass in grams.
Chemists count particles in moles. One mole always contains the same number of particles — Avogadro's constant, 6.02 × 10²³ per mole. The clever part: the mass of one mole in grams is numerically equal to the Mr.
The big numbers in a balanced equation are the mole ratio in which substances react. To find a reacting mass: mass → moles → ratio → moles → mass.
Misconception: mass is conserved — atoms are only rearranged, never created or destroyed. If a product looks lighter (e.g. a gas escapes) the "missing" mass left as gas.
The percentage by mass of an element tells you what fraction of a compound's mass comes from that element — vital for choosing the best fertiliser or ore.
The empirical formula is the simplest whole-number ratio of atoms of each element. The molecular formula shows the actual numbers — always a simple multiple of the empirical formula.
1. Divide each element's mass by its Ar → moles.
2. Divide every answer by the smallest of them.
3. Scale up to whole numbers if needed.
An oxide contains 11.2 g Fe and 4.8 g O.
Fe: 11.2 ÷ 56 = 0.2 | O: 4.8 ÷ 16 = 0.3
Divide by 0.2 → Fe : O = 1 : 1.5 → ×2 → 2 : 3
Empirical formula = Fe₂O₃
Misconception: empirical is not always the molecular formula. Glucose is C₆H₁₂O₆ but its empirical formula is just CH₂O — the simplest ratio.
Water of crystallisation is water chemically bonded into a crystal structure. A salt with it is hydrated; with it removed (by heating to constant mass) it is anhydrous. The dot in a formula shows how many water molecules — the degree of hydration.
1. moles of anhydrous salt = mass ÷ Mr(salt)
2. moles of water lost = mass of water ÷ 18
3. x = moles of water ÷ moles of salt
Tip: Mr of H₂O = 18. Heating "to constant mass" means you keep heating until the mass stops changing — all the water has gone.
The theoretical yield is the maximum mass the equation predicts. The actual yield is what you really collected. The percentage yield compares the two.
Yield is below 100% because some product is lost on separation/transfer, side reactions make other products, or the reaction is reversible and never goes to completion.
Misconception: % yield is not % composition. Yield compares mass got to mass possible for a reaction; composition is the fixed mass make-up of one compound. And neither is atom economy, which is about how much of the product mass is the one you want.
Concentration measures how much solute is dissolved in a given volume. CCEA uses two units: mol/dm³ and g/dm³. Remember 1 dm³ = 1000 cm³, so always convert cm³ → dm³ first.
Misconception: mol/dm³ ≠ g/dm³. One counts particles per litre, the other counts grams per litre. They are only equal if Mr happened to be 1 — which never happens. Convert with the Mr.
A titration finds the volume of one solution that exactly reacts with a measured volume of another. CCEA uses methyl orange or phenolphthalein to spot the end point. You pipette a fixed volume into the flask, then run acid/alkali from the burette until the indicator just changes colour.
At room temperature and pressure (20 °C, 1 atm), one mole of any gas occupies 24 dm³ (24 000 cm³) — its molar volume.
Avogadro's Law: equal volumes of gases at the same temperature and pressure contain the same number of molecules. So gases react in simple volume ratios — you can use the balancing numbers directly on the volumes.
Atom economy measures what fraction of the total mass of products is the desired product. A high atom economy means less waste — important for sustainability and lower costs.
CaCO₃ → CaO + CO₂
Desired CaO: Mr = 40 + 16 = 56
All products: 56 (CaO) + 44 (CO₂) = 100
Atom economy = (56 ÷ 100) × 100 = 56%
Tap a quantity, then tap the box for the unit CCEA uses for it.
Relative formula mass: Mr = sum of all Ar (no units)
Mole: moles = mass ÷ Mr (NA = 6.02×10²³ /mol)
% by mass: (total Ar of element ÷ Mr) × 100
% yield: (actual ÷ theoretical) × 100
Concentration: mol/dm³ = moles ÷ volume(dm³); g/dm³ = mol/dm³ × Mr
Titration: moles = (vol cm³ × conc) ÷ 1000
Gas volume: moles = volume(dm³) ÷ 24 (at RTP)
Atom economy: (Mr desired ÷ total Mr products) × 100
You've covered all of CCEA Quantitative Chemistry — from Mr and the mole to titrations, gas volumes and atom economy. Press Finish to see your score.
You've worked through Quantitative Chemistry for CCEA GCSE Chemistry. 🎉
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Next: test yourself in the Evaluate stage Confidence Quiz, then lock it in with Verify.