Topic 5: Formulae, Equations and Amounts of Substance — Edexcel A-Level Chemistry

Topic 5: Formulae, Equations and Amounts of Substance is the quantitative backbone of A-Level Chemistry. It covers how to represent chemical reactions symbolically, calculate the amounts of reactants and products using moles, and determine empirical and molecular formulae. Mastery of this topic is essential because it underpins stoichiometry, yields, and concentration calculations that appear in every subsequent topic, from energetics to equilibria.

In the Edexcel A-Level specification, this topic introduces the mole concept, Avogadro's constant, molar mass, and the ideal gas equation. You'll learn to balance equations, calculate reacting masses, gas volumes, and solution concentrations, and determine formulae from experimental data. These skills are not only exam-critical but also fundamental for practical work and understanding chemical processes in industry and research.

This topic connects directly to later topics such as energetics (enthalpy changes from bond energies), kinetics (rate calculations), and equilibria (Kc and Kp). A solid grasp here ensures you can handle multi-step calculations and avoid common pitfalls in exams. It's also highly relevant to practical assessments where accurate measurements and calculations are required.

• →The mole is the amount of substance containing 6.02 × 10²³ particles (Avogadro's constant). Molar mass (g mol⁻¹) links mass and moles: moles = mass / molar mass.
• →Empirical formula shows the simplest whole-number ratio of atoms in a compound; molecular formula shows the actual number of atoms. Determine empirical formula from percentage composition or combustion data.
• →Balanced chemical equations must have equal numbers of atoms of each element on both sides. Use state symbols (s, l, g, aq) and ensure coefficients are in the simplest whole-number ratio.
• →Stoichiometry uses mole ratios from balanced equations to calculate reacting masses, volumes of gases (using molar volume 24 dm³ at RTP), and concentrations of solutions (concentration = moles / volume).
• →The ideal gas equation PV = nRT allows calculation of moles, pressure, volume, or temperature for gases under non-standard conditions. R = 8.31 J mol⁻¹ K⁻¹, temperature in Kelvin, pressure in Pa, volume in m³.

• Correct calculation of oxidation numbers in compounds and ions, including peroxides and metal hydrides.
• Correct identification of oxidation and reduction based on electron transfer and oxidation number changes.
• Correct identification of oxidising and reducing agents.
• Correct identification of disproportionation reactions.
• Correct use of Roman numerals to indicate oxidation numbers.
• Correct construction of full ionic equations from ionic half-equations.

• 💡Always check that the sum of oxidation numbers in a neutral compound equals zero and in an ion equals the charge of the ion.
• 💡Remember that oxidising agents are reduced (gain electrons) and reducing agents are oxidised (lose electrons).
• 💡When balancing half-equations, ensure the total charge on both sides is equal.
• 💡Practice identifying oxidation numbers in various contexts, especially for s- and p-block elements.
• 💡Always show your working clearly, including units at each step. Even if your final answer is wrong, you can gain method marks for correct use of formulae like moles = mass/Mr or PV = nRT.
• 💡When calculating empirical formulae from combustion data, remember that all carbon ends up as CO₂ and all hydrogen as H₂O. Calculate moles of C and H from the masses of CO₂ and H₂O, then find the ratio. If oxygen is present, determine its mass by subtraction.
• 💡For percentage yield and atom economy questions, write down the balanced equation first. Percentage yield = (actual yield / theoretical yield) × 100%. Atom economy = (mass of desired product / total mass of reactants) × 100%. Both are often tested together.

• Confusing the direction of electron transfer in oxidation and reduction.
• Incorrectly assigning oxidation numbers in complex ions or species.
• Failing to balance both atoms and charges when constructing ionic half-equations.
• Misidentifying the species being oxidised or reduced in a disproportionation reaction.
• Confusing empirical and molecular formulae: The empirical formula is the simplest ratio, not necessarily the actual formula. For example, ethane has empirical formula CH₃ but molecular formula C₂H₆. Always check if the molar mass matches the empirical formula mass.
• Forgetting to convert units in gas calculations: The ideal gas equation requires volume in m³ (not dm³) and pressure in Pa (not kPa). 1 dm³ = 1×10⁻³ m³, 1 kPa = 1000 Pa. Also, temperature must be in Kelvin (K = °C + 273).
• Using incorrect mole ratios: When calculating reacting masses, ensure you use the stoichiometric coefficients from the balanced equation. For example, in 2H₂ + O₂ → 2H₂O, the mole ratio of H₂ to O₂ is 2:1, not 1:1.