The mole concept and stoichiometry
Work through H1 Chemistry formulae, redox equations and mole calculations, including reacting masses, gas volumes, solutions and titrations.
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Before you begin
Stoichiometry is the arithmetic of chemistry: working out how much of each substance reacts or forms. Almost every calculation follows the same route: convert what you are given into moles, use the ratio from the balanced equation, then convert to what you are asked for. The lessons start with the mole and formulae, then cover reacting masses, solutions and titrations, reacting gas volumes, and redox equations.
Be comfortable with: O-Level chemical calculations, writing chemical equations, atomic structure, and rearranging simple equations.
Learning goals
- Define relative atomic, isotopic, molecular and formula mass
- Define the mole using the Avogadro constant
- Calculate an element's relative atomic mass from the abundances of its isotopes
- Define empirical formula and molecular formula
- Calculate empirical and molecular formulae from combustion data or composition by mass
- Write balanced equations
- Describe and explain redox as electron transfer or a change in oxidation number, using Fe³⁺/Fe²⁺ and MnO₄⁻/Mn²⁺
- Combine half-equations into a redox equation
- Calculate reacting masses from formulae and equations using the mole concept
- Calculate gas volumes in reactions, such as burning hydrocarbons, using the mole concept
- Calculate volumes and concentrations of solutions using the mole concept, to the significant figures asked for
- Deduce stoichiometric relationships from mole calculations
Syllabus statements covered
- define the terms relative atomic, isotopic, molecular and formula mass
- define the term mole in terms of the Avogadro constant
- calculate the relative atomic mass of an element given the relative abundances of its isotopes
- define the terms empirical and molecular formula
- calculate empirical and molecular formulae using combustion data or composition by mass
- write and/or construct balanced equations
- describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state) as exemplified by Fe3+/Fe2+ and MnO4−/Mn2+
- construct redox equations using the relevant half-equations
- perform calculations, including use of the mole concept, involving: — reacting masses (from formulae and equations)
- perform calculations, including use of the mole concept, involving: — volumes of gases (e.g. in the burning of hydrocarbons)
- perform calculations, including use of the mole concept, involving: — volumes and concentrations of solutions [when performing calculations, candidates’ answers should reflect the number of significant figures given or asked for in the question]
- deduce stoichiometric relationships from calculations such as those in (i)
Lessons
Work through them in order.
- Mole and Avogadro constantConvert between mass, amount and number of particles using molar mass and the Avogadro constant.
- Empirical and molecular formulaTurn composition or combustion data into empirical and molecular formulae through a transparent mole-ratio method.
- Reacting masses and limiting reagentUse balanced equations to identify limiting reagents and calculate masses, yields and residual quantities.
- Solution concentration and dilutionUse n = cV and dilution to calculate concentrations, converting volume units correctly.
- Titration calculationsConvert titres into amounts and concentrations using concordant data and equation ratios.
- Reacting gas volumesUse equation coefficients to find reacting gas volumes at the same temperature and pressure.
- Redox and half-equationsTrack electron transfer and oxidation numbers, then construct and combine charge-balanced half-equations.
Practise and check
Or choose
Topic reference
Quick reference
| Question | First move | What to show |
|---|---|---|
| Mass of a product | n = m/M for the substance you know | mole ratio from the equation, then m = nM |
| Limiting reactant | compare moles available with moles needed | use the limiting reactant for every product amount |
| Solution amounts | n = cV, with V in dm³ | divide cm³ by 1000 first |
| Titration | mean of the concordant titres | moles in the titre, mole ratio, then the unknown concentration |
| Gas volumes | at the same temperature and pressure, volume ratio = mole ratio | find the limiting gas; leave out liquids |
| Empirical formula | moles of each element | simplest whole-number ratio |
| Half-equation in acid | balance other atoms, then O with H₂O, H with H⁺, charge with e⁻ | atoms and charge balance |
| Overall redox equation | match the electrons in the two half-equations | add and cancel; no electrons left |
Core knowledge to remember
- Relative atomic mass: the average mass of an atom of an element, relative to one twelfth of the mass of an atom of carbon-12.
- Mole: the amount of substance containing 6.02 × 10²³ particles (the Avogadro constant).
- Empirical formula: the simplest whole-number ratio of atoms of each element in a compound. Molecular formula: the actual number of atoms of each element in one molecule.
- n = m/M, n = cV, and one mole of gas occupies 24.0 dm³ at room temperature and pressure.
- Oxidation: loss of electrons, or an increase in oxidation number. Reduction: gain of electrons, or a decrease in oxidation number.
- An oxidising agent is reduced; a reducing agent is oxidised.
Common mistakes
- Using a 1 : 1 ratio without checking the balanced equation.
- Using a volume in cm³ in n = cV without converting to dm³.
- Using c₁V₁ = c₂V₂ when a reaction takes place. It applies only to dilution.
- Including liquids in a gas volume ratio, or comparing gas volumes measured under different conditions.
- Saying a reducing agent is reduced.
- Leaving electrons in an overall redox equation, or not balancing the charge.
- Rounding in the middle of a calculation, or giving too many significant figures in the answer.