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Q: What does H2 Chemistry Notes: Topic 6 - The Mole Concept and Stoichiometry cover? A: Build systematic problem-solving routines for mole calculations, limiting reagents, redox titrations, and analytical stoichiometry in Core Idea 3 (Mole Concept and Stoichiometry).
Stoichiometry underpins quantitative chemistry-from gas calculations to titration analysis. This note structures the workflow and highlights the must-know techniques for Paper 2 and Paper 3.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13. Core Idea 3 Topic 6 is assessed across Papers 1–3.
Quick revision box
What this topic tests: Mole workflows, limiting reagents, stoichiometric ratios, and titration calculations.
Top mistakes to avoid: Premature rounding; wrong limiting reagent choice; missing units/significant figures.
20-minute sprint plan: 5 min stoichiometry workflow; 10 min limiting/titration practice; 5 min unit + s.f. checks.
Moles from concentration C (mol⋅L−1) and volume V (L).
n=VmV
Moles from gas volume V and molar volume Vm (Data Booklet: 22.7dm3⋅mol−1 at s.t.p.; 24dm3⋅mol−1 at r.t.p.).
Always state units. If conditions differ from the data booklet definitions of s.t.p./r.t.p., revert to the ideal gas equation PV=nRT or use any alternative value provided in the question stem.
2 Stoichiometric Method
Write a balanced equation.
Convert all given quantities to moles.
Use mole ratios from the equation to relate substances.
Convert back to required quantity (mass, volume, concentration).
When numerical work is required, take relative atomic masses and constants directly from the SEAB Chemistry Data Booklet (exams from 2026) rather than rounded memory values.
2.1 Limiting Reagent Logic
Calculate moles of each reactant and compare the ratio with the balanced equation. The reactant yielding the smallest amount of product is limiting. Show working to secure method marks.
A fast check is to compare stoichiometric coefficientn for each reactant: the smaller value identifies the limiting reagent, and all theoretical-yield/purity calculations should then be based on that reagent.
Percentage purity:%purity=total mass of samplemass of pure substance×100
Use mass or moles consistently throughout. For purity problems, the impure mass is often the quantity measured experimentally; set up stoichiometric equations using only the pure component.
4 Redox and Acid-Base Titrations
4.1 Typical Workflow
Write ionic equations (especially for redox).
Convert primary standard volume x concentration into moles.
Apply mole ratio to find moles of analyte.
Convert to requested quantity (concentration, mass, % purity).
4.2 Common Redox Equations
MnOX4X− in acidic medium:
MnOX4X−+8HX++5eX−MnX2++4HX2O
CrX2OX7X2−
SX2OX3X2−
State oxidation numbers to justify electron counts if required.
5 Empirical and Molecular Formulae
Divide percentage or mass data by relative atomic mass to get mole ratio.
Divide all moles by the smallest value to obtain simplest whole-number ratio.
Determine molecular formula using molar mass: Result:MempiricalMmolecular=integer multiplier
Be ready for combustion analysis questions: convert mass of COX2 and HX2O to moles to deduce carbon and hydrogen content.
Carbon dioxide: one-carbon product species used in combustion-analysis mole bookkeeping.
Each COX2 molecule contains one carbon atom, so moles of COX2 directly equal moles of carbon in the original sample.
6 Worked Example
Question:
An impure sample of potassium iodide (KI) weighing 0.700g is titrated with 0.0200molL−1KX2CrX2OX7 in acidic solution. 24.80mL of dichromate is required to reach the endpoint. Determine the percentage purity of the KI sample.
Moles of dichromate:n=CV=(0.0200molL−1)(2.480⋅10−2L)=4.96⋅10−4mol
Moles of iodide: using the 1:6 ratio → n(IX−)=6×4.96⋅10−4mol=2.98⋅10−3mol
Mass of pure KI:m=nM=2.98⋅10−3mol×166.0gmol−1=0.495g
Percentage purity:0.700g0.495g×100=70.7%.
Statement: The KI sample is 70.7% pure.
Remember to report with appropriate significant figures based on experimental data.
7 Practical Tips
Use consistent decimal places in titration tables (e.g. two decimal places for burette readings).
In Paper 4 planning sections, specify standard solutions (e.g. primary standard NaX2COX3 for acid standardisation) and justify choice (stable, high purity).
Mention safety considerations for oxidising agents (KMnOX4, KX2CrX2OX7
8 Common Mistakes
Forgetting dilution effect after mixing solutions.
Applying molar ratios incorrectly when coefficients differ.
Ignoring spectator ions in ionic equations, leading to unbalanced charge.
Using molar volume 24.0dm3⋅mol−1 at non-RTP conditions.
9 Quick Drills
A hydrate CuSOX4⋅xHX2O loses mass from 5.00g to 3.20g upon heating. Determine x.
Calculate the mass of CaCOX3 required to neutralise 25.0mL of 2.00mol⋅L−1
A gaseous hydrocarbon combusts to produce 2.64g of COX2 and 1.08g of HX2O
Check answers with method sheets to ensure your working lines follow the balanced-equation → mole ratio → final quantity structure.
Common exam mistakes
Identifying the wrong limiting reagent: Dividing each reactant's moles by its stoichiometric coefficient gives the correct comparison; the smaller value identifies the limiting reagent. Students who compare raw moles without using coefficients consistently pick the wrong reagent.
Premature rounding of intermediate values: Rounding moles to 2 s.f. mid-calculation introduces cumulative error; carry at least one extra significant figure through each step and round only in the final answer.
Forgetting the dilution effect after mixing solutions: When two solutions are mixed, total volume increases; recalculating concentrations using the new total volume before applying the equilibrium or buffer equation is mandatory.
Using the wrong molar volume for the conditions: 24.0dm3⋅mol−1 applies at r.t.p. and 22.7dm3⋅mol−1 at s.t.p.; using either value at non-standard conditions gives the wrong answer.
Misapplying mole ratios from an unbalanced equation: Ratios must come from the balanced equation; working with an unbalanced equation or using 1:1 ratios by default are both common errors.
Omitting units in stoichiometry working: Missing units (e.g. writing 0.150 instead of 0.150mol) can lead to confusion and risks losing method marks if the examiner cannot follow the logic.
Confusing percentage yield with percentage purity: Percentage yield compares actual to theoretical product amount; percentage purity compares the mass of the desired substance in a sample to the total sample mass - these use different calculation setups.
Frequently asked questions
Do I need to memorise atomic masses for the exam? No. Relative atomic masses are provided in the SEAB Chemistry Data Booklet. However, knowing common values e.g.H=1,C=12,O=16,Na=23,Cl=35.5 speeds up Paper 1 MCQs significantly.
When should I use the molar volume shortcut versus PV = nRT? Use n=V/Vm only when the question states r.t.p. or s.t.p. explicitly and the gas is treated as ideal. In all other cases - non-standard temperatures, pressures, or when ideal-gas assumptions are being tested - use PV=nRT with appropriate unit conversions.
How do I handle a limiting reagent question where one reagent is in excess? Calculate moles of each reactant from the data given, apply the stoichiometric ratio, identify the limiting reagent, and base all subsequent calculations (theoretical yield, percentage yield, remaining excess) on the limiting reagent only.
What is the difference between an empirical formula and a molecular formula? The empirical formula gives the simplest whole-number ratio of atoms; the molecular formula gives the actual number of atoms per molecule. To find the molecular formula, divide the given molar mass by the empirical formula mass to find the integer multiplier.
Struggling with The Mole Concept and Stoichiometry? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
