3.2Stoichiometry

Relative masses of atoms and molecules — Cambridge IGCSE Chemistry 0620 Revision Notes & Topical Questions

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3.2 Relative masses of atoms and molecules — Notes — IGCSE Chemistry notes: Define relative atomic, molecular and formula mass, calculate them, and use the conservation of mass to find reacting mass ratios.Ar is the symbol for relative atomic mass, shown on the atomic symbol in t…

Definition
Common values
Comparing masses
Hydrated salts
Law of conservation of mass
Excess reactant

Section 1 of 7Relative atomic mass

Ar is the symbol for relative atomic mass, shown on the atomic symbol in the Periodic Table.

It is the larger number on the symbol and is always larger than the atomic number, except for hydrogen where they are equal.

3Definition

Definition the average mass of the isotopes of an element compared with one-twelfth of the mass of an atom of carbon-12.

4Standard

Standard atoms are too small to weigh directly, so carbon-12 is the standard with a fixed mass of 12 units and all other atoms are compared with it.

Magnesium has Ar 24, so a magnesium atom is twice as heavy as a carbon atom.

Hydrogen has Ar 1, so a hydrogen atom has one-twelfth of the mass of a carbon-12 atom.

7Units

Units none, because it is a ratio.

8Common values

H 1, C 12, N 14, O 16, Na 23, Mg 24, S 32, Cl 35.5, Ca 40, Cu 63.5.

9Chlorine

Chlorine it has two isotopes, chlorine-35 (75%) and chlorine-37 (25%), so Ar = ÷ 100 = 35.5.
- (75 × 35 + 25 × 37)

10Comparing masses

Ar values are ratios, so 12 g of carbon holds the same number of atoms as 24 g of magnesium.

Section 2 of 7Relative molecular and formula mass

Mr is the sum of the relative atomic masses of all the atoms shown in the formula.
Relative molecular mass is the term for molecules and relative formula mass is the term for ionic compounds.
Method
Write the atoms present from the formula, multiplying through brackets.
Multiply each Ar by the number of atoms of that element.
- Add the results.

Substance	Atoms present	Mr
H₂	2 H	2
H₂O	2 H, 1 O	18
CO₂	1 C, 2 O	44
K₂CO₃	2 K, 1 C, 3 O	138
Ca(OH)₂	1 Ca, 2 O, 2 H	74
(NH₄)2SO₄	2 N, 8 H, 1 S, 4 O	132

Hydrated salts
- CuSO₄.5H₂O has Mr = 63.5 + 32 + 64 + 5 × 18 = 249.5, because the whole 5H₂O is included.

Section 3 of 7Conservation of mass

1Law of conservation of mass

Law of conservation of mass mass cannot be created or destroyed in a chemical reaction, so atoms are rearranged, not made or lost.

In a closed system the total mass of reactants equals the total mass of products.

A balanced equation shows the number of formula units or molecules that take part, so it fixes the mass relationship.

If a gas escapes from an open container the mass appears to fall, but the mass is still conserved in the products.

Section 4 of 7Reacting mass ratios

The ratio of the masses reacting is always the same, whatever the units.

Method

Write the balanced symbol equation.

Work out the formula mass of each substance that is needed.

Multiply each formula mass by its coefficient to get the reacting masses.

Add the reactant masses and the product masses to check that they are equal.

Scale the masses up or down by the same factor to match the mass given.

In 2Ca + O₂ → 2CaO, 80 units of calcium react with 32 units of oxygen to form 112 units of calcium oxide.

80 g, 40 kg or any other unit keeps the same ratio, so doubling the methane in CH₄ + 2O₂ → CO₂ + 2H₂O doubles the carbon dioxide.

10Excess reactant

Excess reactant when one reactant is in excess, all of the other reacts, so the other decides the product mass.

Section 5 of 7More from the specification

1Carbon-12 standard and atomic mass

Carbon-12 standard and atomic mass comparing atomic isotopic masses against unified carbon-12 standards provides a relative means to determine molecular weights.

2Determining Empirical Formula from Percentage Composition

The empirical formula represents the simplest whole-number ratio of atoms of each element in a compound, calculated by dividing mass percentages by relative atomic masses.

3Precipitation Reactions and Ionic Balance

When aqueous solutions of copper(II) sulfate and sodium hydroxide mix, blue insoluble copper(II) hydroxide precipitates while soluble sodium sulfate remains in solution.

4Hydrated Salts and Water of Crystallisation

Hydrated copper(II) sulfate crystals contain chemically bonded water of crystallisation (CuSO₄·5H₂O),
which is driven off upon heating to leave white anhydrous copper(II) sulfate powder.

5Charge

Q = It, measured in coulombs.
What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

6Sulfate

Salt of sulfuric acid.

7Hydroxide

A compound containing the OH⁻ ion, formed by Group I metals with water.

8Nitrate

Salt of nitric acid.

9Empirical formula

The simplest whole-number ratio of atoms or ions.

10Sodium chloride

Compound of sodium and chlorine; common salt.

Section 6 of 7Key terms

Term	Meaning
Relative atomic mass	the average mass of an element's isotopes compared with 1/12 of the mass of a carbon-12 atom.
Relative molecular mass	the sum of the relative atomic masses of all atoms in a molecule.
Relative formula mass	the sum of the relative atomic masses in the formula of an ionic compound.
Conservation of mass	the total mass of the reactants equals the total mass of the products.
Coefficient	the number in front of a formula in a balanced equation.

Section 7 of 7Exam technique

Giving units for Ar or Mr — relative masses have no units

Ignoring the brackets or the water of crystallisation when adding up a formula — multiply every atom inside brackets by the number outside and include the whole 5H₂O in CuSO4.5H2O

Using the atomic number instead of Ar — Ar is the larger number

Forgetting the coefficient or not balancing first — multiply Mr by the number in front of the formula, so 2O₂ is 64, not 32

Rounding the Ar of chlorine to 35 — use 35.5 unless told otherwise

Last updated: 18 October 2026