Quantitative Chemistry: Conservation of Mass, the Mole and Molar Calculations

This guide covers Quantitative chemistry, one of the ten topic areas of OxfordAQA International GCSE Chemistry 9202: conservation of mass and the quantitative interpretation of chemical equations, amount of substance in relation to masses of pure substances, the mole concept, and using molar concentrations of solutions and amount of substance in relation to volumes of gases.

Before studying this

This topic assumes you’re already confident writing and balancing chemical equations (covered under Chemical changes) and comfortable with basic proportion and ratio arithmetic. Quantitative chemistry is where those skills turn into real calculations — working out exactly how much of each substance reacts and forms, which underpins practical chemistry from titrations to industrial-scale reactions.

Conservation of mass

No atoms are created or destroyed in a chemical reaction — they’re only rearranged. This means the total mass of the reactants always equals the total mass of the products, provided nothing enters or leaves the reaction system as a gas.

A balanced chemical equation is a quantitative statement, not just a description: the numbers in front of each formula (the stoichiometric coefficients) tell you the exact ratio in which particles react and form. For example, in

2Mg + O₂ → 2MgO

two moles of magnesium react with one mole of oxygen to form two moles of magnesium oxide — a ratio you can scale up or down to any actual quantity.

The mole concept

The mole (mol) is the unit chemists use to count particles, because atoms and molecules are far too small to count individually. One mole of any substance contains the same number of particles — the Avogadro constant, approximately 6.02 × 10²³ per mole.

The molar mass of a substance (units g mol⁻¹) is the mass of one mole of it, numerically equal to its relative formula mass. This gives the key relationship used throughout this topic:

amount (mol) = mass (g) ÷ molar mass (g mol⁻¹)

Reacting-mass calculations

Combining conservation of mass with the mole lets you calculate exactly how much product forms from a given mass of reactant, or how much reactant is needed to form a given mass of product.

Worked example. What mass of magnesium oxide forms when 6.0 g of magnesium burns completely in oxygen? (Aᵣ: Mg = 24, O = 16)

1. Moles of Mg = 6.0 ÷ 24 = 0.25 mol
2. From 2Mg + O₂ → 2MgO, the ratio of Mg to MgO is 1:1, so moles of MgO formed = 0.25 mol
3. Mass of MgO = moles × molar mass = 0.25 × 40 = 10.0 g

Molar concentration of solutions

The concentration of a solution is usually expressed in moles per cubic decimetre (mol dm⁻³):

amount (mol) = concentration (mol dm⁻³) × volume (dm³)

remembering that 1 dm³ = 1000 cm³, so a volume in cm³ must be divided by 1000 before use in this equation.

Worked example. How many moles of hydrochloric acid are in 25.0 cm³ of a 2.00 mol dm⁻³ solution?

amount = concentration × volume = 2.00 × (25.0 ÷ 1000) = 0.0500 mol

Not every titration is a simple 1:1 reaction, and the mole ratio must always come from the balanced equation. For example, in H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O, one mole of sulfuric acid reacts with two moles of sodium hydroxide. If 25.0 cm³ of 0.0500 mol dm⁻³ sulfuric acid (1.25 × 10⁻³ mol) is exactly neutralised by 20.0 cm³ of sodium hydroxide solution, the moles of NaOH used are 2 × 1.25 × 10⁻³ = 2.50 × 10⁻³ mol, giving a concentration of (2.50 × 10⁻³ × 1000) ÷ 20.0 = 0.125 mol dm⁻³ — assuming a 1:1 ratio here would halve this answer.

Amount of substance and gas volumes

At room temperature and pressure, one mole of any gas occupies the same volume — conventionally taken as 24 dm³ mol⁻¹ (24,000 cm³ mol⁻¹) at GCSE level:

amount (mol) = volume of gas (dm³) ÷ 24

This lets you calculate the volume of gas produced or consumed in a reaction directly from the balanced equation, using exactly the same mole-ratio method as a reacting-mass calculation.

Worked example. What volume of carbon dioxide, at room temperature and pressure, is produced when 2.50 g of calcium carbonate reacts completely with excess dilute hydrochloric acid? CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ (Aᵣ: Ca = 40, C = 12, O = 16)

1. Molar mass of CaCO₃ = 40 + 12 + 3(16) = 100 g mol⁻¹
2. Moles of CaCO₃ = 2.50 ÷ 100 = 0.0250 mol
3. The ratio CaCO₃ : CO₂ is 1 : 1, so moles of CO₂ formed = 0.0250 mol
4. Volume of CO₂ = 0.0250 × 24 = 0.600 dm³

Common mistakes

• Forgetting to convert cm³ to dm³ before a concentration calculation. Concentration in mol dm⁻³ always needs the volume in dm³ (divide cm³ by 1000), not cm³ directly.
• Reading the mole ratio off the formulae instead of the balancing numbers. The stoichiometric coefficients in the balanced equation, not the subscripts inside a formula, give the reacting ratio between separate substances.
• Assuming mass is conserved even when a gas escapes an open container. Conservation of mass applies to the whole reacting system — if a gas is allowed to escape (an open flask, for example), the measured mass of the remaining solid or solution will appear to decrease, even though no atoms have actually been destroyed.
• Mixing up molar mass and molar gas volume. Molar mass converts between mass and moles; molar gas volume converts between gas volume and moles — they are used in different steps and are not interchangeable.

Quick revision checklist

• Conservation of mass and its link to the stoichiometric coefficients in a balanced equation
• The mole as a counting unit, and the Avogadro constant
• amount = mass ÷ molar mass
• amount = concentration × volume (in dm³) for solutions
• amount = gas volume (dm³) ÷ 24 at room temperature and pressure

Related resources

• Formulae, Equations and the Mole (Cambridge IGCSE / O Level) — the equivalent topic from a different awarding body
• Amount of Substance (OCR A Level) — how these same mole-calculation skills extend at A Level
• OxfordAQA International GCSE Chemistry hub

Written against OxfordAQA International GCSE Chemistry 9202, specification updated November 2022, https://www.oxfordaqa.com/qualifications/international-gcse-chemistry/, verified 2026-08-18. Always check the current specification for your examination year.