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's current H2 Chemistry (9476) syllabus PDF is labelled for 2026, and the current Chemistry Data Booklet is labelled 8873/9476/9813 for use from 2026 in non-practical papers. Core Idea 3 Topic 6 is assessed across Papers 1-3.
The core idea is simple: Stoichiometry is a conversion workflow: given quantity to moles, mole ratio, then required quantity.
Use it as a working check: Write the balanced equation before calculating. The coefficients decide the ratio, not the numbers that appear in the question first.
Then go one layer deeper: Example: if 2mol of A reacts with 1mol
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Attempt the matching topic bank before reopening the notes. Use each missed idea to decide what to review next.
; the smaller value identifies the limiting reagent.
Route map: choose the stoichiometry pathway first
If the question gives you...
Start with...
Then connect to...
Trap to avoid
Mass and molar mass
Convert mass to moles
Balanced equation ratio, then required mass or amount
Do not compare masses directly when coefficients differ.
Concentration and volume
Convert volume to litres, then use moles equals concentration times volume
Titration ratio, dilution, or concentration of the unknown
Do not leave millilitres inside a concentration in mol per litre calculation.
Gas volume at r.t.p. or s.t.p.
Use the stated molar volume
Mole ratio, then gas volume, mass, or concentration
Do not use molar volume when the question gives non-standard temperature or pressure.
Two reactants with amounts
Divide each mole amount by its coefficient
The smaller normalised value gives the limiting reagent
Do not choose the reactant with fewer raw moles by default.
Actual yield or impure sample data
Decide whether the question asks yield or purity
Theoretical amount for yield, pure component amount for purity
Do not swap the denominator in percentage yield and percentage purity.
Empirical formula or combustion data
Convert each element source to moles
Simplest whole-number ratio, then molecular multiplier if given
Do not round ratios too early before checking for simple fractions.
Use this map before calculating. Most errors in Topic 6 happen before the arithmetic, when the wrong pathway is chosen.
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.
1 Fundamental Relationships
Formula
Description
n=Mm
Moles from mass and molar mass.
n=CV
Moles from concentration C (mol⋅L−1) and volume V (L).
n=VmV
Moles from gas volume V and molar volume Vm
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 Aliquot and dilution checkpoint
Before using a titre, identify which solution volume the titre actually reacts with. This keeps stock-solution concentration, diluted-solution concentration, aliquot volume, and average titre from being swapped.
Quantity in the question
What it represents
First calculation move
Pipetted aliquot
Fixed volume transferred into the conical flask
Use this as the analyte volume in the mole ratio, not the full volumetric-flask volume.
Average titre
Volume delivered from the burette to react with the aliquot
Convert to litres, then use n=CV for the titrant.
Volumetric-flask volume
Final volume after dilution
Use it only when scaling from aliquot concentration back to the diluted solution or stock solution.
Dilution statement
How the stock solution was made less concentrated
Apply the dilution factor after finding the concentration of the diluted solution.
Misconception check: the titre does not usually react with the whole volumetric flask. It reacts with the aliquot in the flask, so scale back to the original solution only after the mole ratio step is complete.
4.3 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.
Combustion formula bookkeeping checkpoint
Combustion product or data
Mole link
Common trap
Mass of COX2
n(C)=n(COX2)
Multiplying carbon moles by 2 because carbon dioxide has two oxygen atoms
Mass of HX2O
n(H)=2n(HX2O)
Original sample contains only C, H, and O
Find oxygen by mass difference first: m(O)=m(sample)−m(C)−m(H)
Subtracting mole amounts instead of masses
Molar mass is given
Compare molecular formula mass with empirical formula mass
Rounding the empirical mole ratio before checking simple fractions
Worked check: an organic compound contains only C, H, and O. A 0.300g sample gives 0.440g of COX2 and 0.180g of HX2O.
n(C)=n(COX2)=44.00.440=0.0100mol
n(H)=2n(HX2O)=2(18.00.180)=0.0200mol
Masses: m(C)=0.0100×12.0=0.120g, m(H)=0.0200×1.0=0.0200g, so m(O)=0.300−0.120−0.0200=0.160g and n(O)=0.160/16.0=0.0100mol. The empirical formula is CHX2O.
Misconception check: oxygen by difference is a mass step before it is a mole-ratio step.
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
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.
