JC2 Chemistry Tutor Hougang | How do you calculate an unknown Fe²⁺ concentration from an acidified manganate(VII) titration? A student can read a burette confidently and still lose marks by using a 1:1 mole ratio where five electrons are actually transferred. The critical step is the balanced redox equation: one MnO₄⁻ ion accepts five electrons as it is reduced to Mn²⁺ in acid, while each Fe²⁺ ion loses one electron to become Fe³⁺. Concentration follows only after the ratio is correct.
For H2 A-Level Chemistry 9476, manganese and chromium redox systems, balanced half-equations, volumetric analysis and practical data interpretation are connected JC2 skills. The 2027 SEAB syllabus explicitly includes Fe³⁺/Fe²⁺, MnO₄⁻/Mn²⁺ and Cr₂O₇²⁻/Cr³⁺ as redox systems. This guide uses original Fe²⁺ and oxalate calculations, discusses endpoints and measurement uncertainty, and shows how a tutor can distinguish a mole-ratio mistake from an experimental error.
All laboratory reagents and redox titrations require suitable supervised facilities and safe disposal. This article is for written chemical reasoning and interpreting supplied experimental data, not independent home experimentation. eduKateSG’s small groups of up to three students meet at 8 Fourth Avenue near Sixth Avenue MRT, not in Hougang; confirm lesson arrangements and commute before enrolling.
The 2027 H2 Chemistry assessment scope
The 2027 SEAB H2 9476 syllabus includes relevant redox systems of manganese(VII), chromium(VI) and iron(II)/(III), and practical assessments involving volumetric work, observations, data analysis, accuracy and precision. Students need to balance oxidation and reduction, use the mole ratio in quantitative calculations and interpret experimental evidence. A successful solution depends on both the equation and the quantities measured.
Do not assume every past-paper experiment has the same reagents or the same endpoint colour. The detailed practical instructions on the actual paper control what is performed. The theory here equips students to interpret a supplied Fe²⁺ or oxalate titration without pretending the article replaces laboratory competence.
Oxidation and reduction are electron transfers
Oxidation involves loss of electrons and reduction involves gain of electrons. Fe²⁺→Fe³⁺+e⁻ is oxidation, while MnO₄⁻+8H⁺+5e⁻→Mn²⁺+4H₂O is reduction under the appropriate acidic conditions. The oxidising agent accepts electrons and is itself reduced; here manganate(VII) is the oxidising reagent in the standard school example.
A student who calls Fe²⁺ the oxidising agent because iron appears in the name has ignored the actual charge change. Identify the element’s oxidation number before naming the agent. The direction is determined by the half-equations rather than a mnemonic about which reagent was added first.
Why acid appears in the manganate half-equation
The reduction of MnO₄⁻ to Mn²⁺ under acidic conditions requires protons and water to balance hydrogen and oxygen. The full half-equation is MnO₄⁻+8H⁺+5e⁻→Mn²⁺+4H₂O. A student who omits H⁺ and H₂O is not describing the same balanced redox system, and the stated acidic environment matters to the expected manganese product.
Check atoms: one Mn, four O on both sides, and eight H after including four water molecules. Check charge: left is −1+8−5=+2, matching the Mn²⁺ on the right. This is a clean method for confirming a half-equation without relying only on memory.
Iron(II) oxidation is a one-electron process
Fe²⁺→Fe³⁺+e⁻ is the relevant iron oxidation half-equation in the standard example. The oxidation-state change is +2 to +3, so one electron is released per iron ion. Because manganate(VII) takes five electrons, five Fe²⁺ ions must be oxidised for every MnO₄⁻ ion reduced in the balanced reaction.
Students often skip this electron-matching step and assume one mole of titrant reacts with one mole of analyte. The actual titration stoichiometry is not determined by the number of reagent names but by the electrons exchanged and the balanced equation.
The full acidified MnO4− and Fe2+ equation
Combining the half-equations gives MnO₄⁻+5Fe²⁺+8H⁺→Mn²⁺+5Fe³⁺+4H₂O. This is the central ratio for the standard Fe²⁺–manganate(VII) calculation in acidic medium. One mole of MnO₄⁻ corresponds to five moles of Fe²⁺ consumed, assuming the reaction completes as stated and no competing processes affect the result.
Read the total charge as a final check: the left is −1+10+8=+17 and the right is +2+15=+17. The oxygen and hydrogen also balance. A student who has a chemically balanced net ionic equation can use its coefficients with confidence.
A titre is a delivered volume, not a final burette reading
If a burette starts at 0.35 cm³ and ends at 24.95 cm³, the delivered volume is 24.60 cm³. That difference is the titre. A student who copies the final reading as the titre has used the wrong measured quantity, even when the later mole calculation is otherwise correct. In some apparatus designs or starting conditions, the initial reading may not be zero.
Use a table with initial reading, final reading and calculated titre. Respect the apparatus’s recording convention and avoid adding spurious decimal places. A clean volume calculation is the first quantitative control before translating titre into moles.
Worked Fe²⁺ concentration from a titre
Suppose a 25.0 cm³ Fe²⁺ sample requires 20.00 cm³ of 0.0200 mol dm−3 acidified manganate(VII) under the specified balanced reaction. The manganate amount is 0.0200×0.02000=4.00×10−4 mol. The Fe²⁺ amount is five times that, 2.00×10−3 mol. Dividing by 0.0250 dm³ gives Fe²⁺ concentration 0.0800 mol dm−3.
The calculation has three steps: convert delivered volume to dm³, find oxidant moles, then use the five-to-one ratio before dividing by sample volume. If the answer is 0.0160 M, the student probably used the wrong stoichiometric factor. If it is 80.0 M, the conversion from cm³ may have been missed.
Why total flask volume is not always aliquot volume
In an ordinary titration calculation, the concentration sought for the aliquot is found from the moles reacted divided by that measured aliquot volume. The titrant volume added is used to determine titrant moles, not automatically added to the aliquot volume for the requested original analyte concentration. The combined volume can matter for a different question about the final solution’s concentration.
Students should identify whether they are calculating the original sample concentration, the post-reaction mixture or a species at equilibrium. Using the wrong denominator is a physical misunderstanding, not just arithmetic.
Oxalate loses two electrons per ion
An oxalate ion C₂O₄²⁻ can be oxidised to carbon dioxide through the half-equation C₂O₄²⁻→2CO₂+2e⁻. Two carbon atoms appear on each side, four oxygens appear in the two CO₂ molecules, and charge is −2 on both sides when electrons are included. The reaction is a two-electron oxidation per oxalate ion in the model.
This differs from iron(II), which releases one electron per ion. Therefore the mole ratio with manganate(VII) cannot be copied from the iron titration. Balance electrons afresh whenever the analyte changes.
Oxalate and manganate(VII) balance as 5:2
Combining the two-electron oxalate oxidation with the five-electron manganate reduction requires ten electrons overall. Multiply the manganate half-equation by two and the oxalate half-equation by five. The resulting acidic net ionic equation is 2MnO₄⁻+5C₂O₄²⁻+16H⁺→2Mn²⁺+10CO₂+8H₂O.
A student who writes one mole of manganate for one mole of oxalate has ignored both half-equations. The coefficients two and five are not optional: they express the ten-electron balance. Check total charge and all atoms before beginning any concentration calculation.
Worked oxalate titre problem
Suppose a 25.0 cm³ sample contains 0.0400 mol dm−3 oxalate. Its amount is 0.0250×0.0400=1.00×10−3 mol oxalate. From the net equation, the required manganate amount is (2/5)×1.00×10−3=4.00×10−4 mol. If manganate concentration is 0.0200 mol dm−3, the corresponding theoretical titre is 0.0200 dm³, or 20.0 cm³.
Notice that the numerical titre matches the earlier Fe²⁺ example, but for an entirely different analyte amount and ratio. Reaching the same number is not evidence the chemistry is the same. The balanced equations explain why both calculations happen to yield that value.
The endpoint is an observation, not the mole ratio
In an ordinary acidified manganate(VII) titration, the reagent’s intense purple colour can act as its own visual indicator, with a faint persistent pink colour used as an endpoint cue under standard laboratory protocols. The observed endpoint is an experimental measurement, whereas the stoichiometric equivalence point is defined by the balanced reaction. The two are intended to be close, but they are not identical by definition.
A student should record the actual observation specified in the paper rather than invent an ideal colour change from memory. Interpreting an endpoint requires understanding reagent and analyte, the reaction’s progress and the need for appropriately controlled practical technique.
Why acid choice can matter to a redox system
The acidic medium is a chemical reactant in the MnO₄⁻ reduction half-equation, but not every acid supplies only harmless spectator ions under all circumstances. In some strongly oxidising titrations, chloride ions from hydrochloric acid can be susceptible to side oxidation, while other acids may introduce their own oxidising chemistry. The experimental instructions must therefore specify suitable conditions and reagents.
This is a conceptual reason for following a validated supervised protocol, not a home procedure. For examination theory, recognise when a proposed reagent can interfere chemically instead of assuming any liquid labelled acid is interchangeable.
Why oxalate reaction kinetics need context
Oxalate–manganate(VII) redox chemistry can show rate behaviour that depends on temperature and on manganese-containing species formed during the reaction. A question may supply observations or a graph and ask the student to explain changing reaction rates. That is a kinetics issue layered on top of a thermodynamically possible and stoichiometrically balanced redox process.
Do not infer an exact practical temperature or procedure from this general statement. The syllabus assessment may provide the relevant conditions. A tutor’s paper-based task can still ask why balanced overall stoichiometry does not automatically determine how quickly an endpoint is reached.
Dichromate(VI) provides a six-electron comparison
In acidic solution the dichromate(VI) reduction can be written Cr₂O₇²⁻+14H⁺+6e⁻→2Cr³⁺+7H₂O. Coupled with Fe²⁺→Fe³⁺+e⁻, one mole of dichromate corresponds to six moles of Fe²⁺ under the balanced reaction. This is a different stoichiometric relationship from the five-to-one manganate(VII) comparison.
A student who memorises the Fe²⁺ titration answer but cannot explain the six-electron dichromate half-equation is vulnerable to a changed reagent. The correct ratio comes from the number of electrons transferred in the appropriate redox system.
The full dichromate–iron reaction
Combining the acidic dichromate half-equation with six Fe²⁺ oxidation steps gives Cr₂O₇²⁻+6Fe²⁺+14H⁺→2Cr³⁺+6Fe³⁺+7H₂O. The charge on the left is −2+12+14=+24, while the right is +6+18=+24, so the ionic charge balances. The atoms also balance.
This is an excellent practice comparison with manganate(VII), because the student must choose whether to multiply Fe²⁺ by five or six based on the named oxidising reagent, not from a generic titration rule.
Concordant titres do not guarantee accuracy
Several nearly identical titres demonstrate repeatability and may indicate good precision, but a systematic error can affect all of them. A reagent solution prepared with incorrect concentration or a consistently misread instrument can produce concordant yet inaccurate results. The correct evaluation should distinguish random scatter from a persistent bias.
If a Paper 4 data-analysis question asks for improvement, explain how the proposed change addresses the specific error. Repeating an experiment can help estimate random variability; it does not automatically remove a known systematic concentration error.
The uncertainty of a titre involves two readings
A titre is the difference between initial and final burette readings, so both readings contribute to its measurement uncertainty. The exact absolute uncertainty should follow the apparatus specification and the task’s recording rules. A student who blindly quotes the uncertainty for one reading as the entire delivered-volume uncertainty may understate the measurement limitation.
In a written uncertainty calculation, use the stated instrument information instead of making up a universal burette value. Convert the resulting absolute uncertainty into a percentage of the measured titre only when the question requests that comparison.
Percentage uncertainty depends on the size of the titre
For the same absolute volume-reading uncertainty, a small titre has a larger percentage uncertainty than a larger titre. This is a mathematical consequence of percentage uncertainty being absolute uncertainty divided by the measured quantity times 100. That can inform the interpretation of experimental precision and sensible reagent concentration choices under a validated laboratory task.
However, a large titre is not automatically better under every condition: apparatus capacity, time, endpoint clarity and safety also matter. A reasoned experimental recommendation should refer to the actual measurement constraints.
Do not confuse titre precision with stoichiometric validity
A perfectly reproducible volume does not establish a correct analyte concentration if the balanced equation is wrong. For example, using a one-to-one Fe²⁺/MnO₄⁻ ratio in all calculations would give a systematic factor-of-five error even with identical measured titres. The analytical model is just as important as the measurement.
Ask the student to derive the electron ratio independently before using an average titre. This separates chemical reasoning from practical data quality, two different sources of marks in H2 Chemistry.
What an incomplete reaction changes
The ideal calculation assumes that the chosen redox reaction is effectively complete at the relevant endpoint and that other species do not consume significant titrant. If a sample contains another reducing agent, its reaction could also affect the measured volume. An experimental result should therefore be interpreted within the sample composition and chemical selectivity described by the question.
Do not invent an exact concentration from a titre if the source information makes the analyte ambiguous. A strong evaluation may explain what further selectivity evidence or control would be needed.
One ratio check before the calculator
For Fe²⁺ with MnO₄⁻ in acid, ask: how many electrons does manganate accept? Five. How many does one Fe²⁺ donate? One. Therefore the ratio Fe²⁺:MnO₄⁻ is 5:1. For oxalate, each ion donates two electrons, so five oxalates balance two manganates. This mental check takes seconds and can prevent the most expensive exam errors.
Then write the full equation and only then insert titre data. A tutor should mark which stage failed, rather than call every wrong final concentration careless.
A four-column titration calculation layout
Keep the working in a visible order: measured titre and its units, titrant moles, analyte moles using the balanced redox ratio, and the required analyte concentration or mass. If the sample was diluted earlier, account for the corresponding dilution factor in an additional clearly labelled step. Do not mix the original stock volume with the aliquot without stating which solution each measurement describes.
This layout makes it easier to audit a numerical answer and easier for the student to discover the first wrong step. It can be used with other volumetric analysis reactions provided their actual chemical equations are balanced.
Safety is not optional
Acidified strong oxidants, metal solutions and oxalate reagents require appropriate safety equipment, authorised laboratory protocols and proper waste management. They should not be used for unsupervised home experimentation. A paper guide can teach half-equations, stoichiometry, observations and uncertainty without serving as a procedure for handling hazardous chemicals.
Students preparing for Paper 4 should obtain supervised practical training through suitable educational facilities. A desk-based tutor should be honest about the difference between explaining a laboratory technique and providing hands-on practical instruction.
Case 1 — One MnO₄⁻ does not mean one Fe²⁺
Question: A student writes the net reaction as MnO4−+Fe2+→Mn2++Fe3+ under acidic conditions.
Reasoning: The equation is not electron-balanced. Manganate(VII) accepts five electrons to become Mn²⁺, while one Fe²⁺ ion releases only one. Five iron(II) ions must therefore be oxidised for each manganate ion reduced. Including acid and water gives the correctly balanced acidic ionic equation. A wrong ratio propagates into every subsequent concentration calculation.
Changed check: Replace Fe²⁺ with a two-electron reducing species and derive a fresh ratio.
Case 2 — Check charges on a half-equation
Question: The reduction half-equation is written without electrons, H+ or water.
Reasoning: For MnO4−→Mn2+ in acid, the complete equation is MnO4−+8H++5e−→Mn2++4H2O. Atoms and total charge must balance. The left charge is +2 after including electrons and protons; the right charge is also +2. This gives a defensible electron count rather than a remembered numeral.
Changed check: Try writing the dichromate half-equation and verify its charge balance independently.
Case 3 — A final burette number is not the titre
Question: A burette begins at 1.20 cm3 and finishes at 22.75 cm3.
Reasoning: The delivered titre is 22.75−1.20=21.55 cm3. The final reading alone is not the volume added. A calculation using 22.75 cm3 would overstate titrant moles and could give a plausible-looking concentration with a systematic error. Always keep initial, final and delivered volume in separate columns.
Changed check: Give a run starting at a non-zero reading close to 10 cm3 and retest the subtraction.
Case 4 — Convert centimetres cubed to decimetres cubed
Question: A titre of 20.00 cm3 is used with titrant concentration expressed in mol dm−3.
Reasoning: Convert 20.00 cm3 to 0.02000 dm3 before multiplying by concentration. Multiplying 20.00 directly by 0.0200 mol dm−3 introduces a factor of one thousand. The correct titrant amount is 4.00×10−4 mol, not 0.400 mol. Unit conversion is a physical meaning check, not a calculator preference.
Changed check: Change the titre to 18.65 cm3 and calculate the titrant moles with appropriate precision.
Case 5 — Calculate Fe²⁺ amount from manganate amount
Question: An acidified manganate titre corresponds to 4.00×10−4 mol MnO4−.
Reasoning: The balanced acid reaction consumes five Fe²⁺ per MnO4−, so the Fe²⁺ amount in the aliquot is 2.00×10−3 mol. The moles are not equal, and the factor five should be justified from electron balancing rather than copied from a prior worked answer. The step follows the chemistry independently of whether the sample volume is 10 or 25 cm3.
Changed check: Use a different oxidant with six-electron reduction and ask how the ratio changes.
Case 6 — Identify the correct sample volume
Question: A 25.0 cm3 Fe²⁺ aliquot is titrated with 20.00 cm3 manganate solution.
Reasoning: The original Fe²⁺ concentration is its reacted amount divided by 0.0250 dm3, not by 0.0450 dm3 of combined liquid. The combined volume might matter for a different final-mixture concentration question, but it is not the original aliquot denominator. Distinguish which solution’s concentration the examiner requests.
Changed check: Repeat with a 10.0 cm3 aliquot and show why the same reacted moles imply a different original concentration.
Case 7 — A diluted stock solution requires another factor
Question: A measured aliquot came from a stock that was diluted tenfold before titration.
Reasoning: The titration determines concentration in the diluted solution being analysed. To report the original stock concentration, account for the stated dilution factor after the titration calculation. Simply copying the diluted aliquot concentration as the original stock would be wrong. The correct approach follows moles through each preparation stage and avoids treating the words ‘stock’ and ‘aliquot’ as synonymous.
Changed check: Give a two-stage dilution and ask which volumes belong to the sample-preparation ledger.
Case 8 — Oxalate is a two-electron donor
Question: The same manganate reagent is now used for an oxalate sample rather than Fe²⁺.
Reasoning: Oxalate C2O4²−→2CO2+2e− releases two electrons per ion, whereas Fe²⁺→Fe³⁺+e− releases one. This changes the stoichiometric ratio. The acidic overall reaction uses two MnO4− for five oxalate ions, balancing ten electrons on each side. Do not reuse the five-to-one Fe²⁺ ratio.
Changed check: Calculate the amount of oxalate corresponding to a specified amount of MnO4− using the 5:2 ratio.
Case 9 — Oxalate amount from a theoretical titre
Question: A 0.0200 M manganate titrant delivers 20.0 cm3 to a stoichiometric oxalate reaction.
Reasoning: Titrant amount is 0.0200×0.0200=4.00×10−4 mol. Because 2 MnO4− react with 5 oxalate ions, the oxalate amount is (5/2)×4.00×10−4=1.00×10−3 mol. Dividing by a 25.0 cm3 aliquot gives oxalate concentration 0.0400 M under the stated assumptions.
Changed check: Change the oxidant concentration and ask for the new expected titre while keeping analyte moles fixed.
Case 10 — The dichromate ratio is six, not five
Question: A question uses Cr2O7²− reduction to Cr3+ with Fe²⁺ as the reducing species.
Reasoning: The acidic dichromate half-equation accepts six electrons, so one mole of dichromate corresponds to six moles of Fe²⁺. Reusing the manganate five-electron factor would be a reagent-identity mistake. The half-equation controls stoichiometry more reliably than the fact that both solutions are oxidising agents.
Changed check: Write the balanced dichromate–iron net equation and compare its coefficients with the manganate version.
Case 11 — Acid is not merely a solvent label
Question: The manganese reduction is described as occurring in acidic medium.
Reasoning: The hydrogen ions appear explicitly in the acidic half-equation and help balance oxygen and hydrogen as MnO4− becomes Mn2+. The chosen medium can affect product chemistry and possible side reactions. A chemically appropriate answer should not omit the stated acidity or assume the same manganese product arises under every possible pH condition.
Changed check: Compare the question’s stated acid environment with a hypothetical alkaline environment and identify why a new model might be needed.
Case 12 — A colour endpoint is not equivalence by definition
Question: A hypothetical practical describes a persistent faint pink endpoint for an acidified manganate titration.
Reasoning: That visual cue marks the observed endpoint under the specified procedure. The equivalence point, by contrast, is the exact stoichiometric condition set by the balanced redox equation. A good method aims for a close relationship between them, but measurement uncertainty and practical conditions can affect the observed volume. Do not treat the two definitions as identical.
Changed check: Explain which part of the titration calculation uses measured endpoint volume and which uses chemical stoichiometry.
Case 13 — Concordance and accuracy are different
Question: Three recorded titres agree to within a small range.
Reasoning: They show consistency and may support precision, but do not prove the titrant concentration was correct or that the endpoint was free of systematic bias. Repeating a consistently flawed procedure can give very similar wrong measurements. An appropriate evaluation names the specific limitation and recommends a remedy that addresses it.
Changed check: Given a fictional miscalibrated instrument, decide whether another repeat alone would fix the bias.
Case 14 — Percentage uncertainty needs a denominator
Question: An experiment reports an absolute titre uncertainty but asks for percentage uncertainty.
Reasoning: Use percentage uncertainty=(absolute uncertainty of the titre)/(measured titre)×100 under the specified treatment. A smaller titre gives a larger percentage uncertainty for the same absolute reading limitation. The exact absolute figure must come from apparatus information or examination instructions rather than an invented universal number.
Changed check: Compare uncertainty percentages for 10.0 cm3 and 25.0 cm3 titres under the same assumed absolute uncertainty.
Case 15 — Heat or catalyst can affect rate, not stoichiometry
Question: A redox reaction is described as slow under one set of conditions but faster under another.
Reasoning: The balanced overall equation still governs electron and mole ratios for the stated products, even if the reaction rate changes. A catalyst or temperature effect does not authorise changing the electron count merely to match an observed endpoint time. Separate kinetics from stoichiometric equivalence when interpreting practical data.
Changed check: Compare a fast and slow hypothetical run with the same total reacted amounts and explain why theoretical titrant moles match.
Case 16 — The same titre may mean different analyte concentrations
Question: Two experiments use identical titres of manganate but different analyte aliquot volumes.
Reasoning: The titrant moles and, for the same redox chemistry, analyte moles may match, but original analyte concentrations differ because amount is divided by each aliquot volume. A student should not declare that equal titre always means equal concentration. The experiment’s chosen sample volume is essential.
Changed check: Use 10.0 cm3 and 25.0 cm3 aliquots with the same Fe²⁺ moles and compare concentrations.
Case 17 — Don’t confuse iron colours with the electron ratio
Question: An observation reports a change in the colour of an iron-containing solution after oxidation.
Reasoning: Colour may support a chemical interpretation but does not itself specify the electron-transfer coefficient. Fe²⁺→Fe³⁺ loses one electron, established through oxidation states and balanced charge. Different complex environments can change colours without the same oxidation-state change. Use the equation and data rather than infer stoichiometry solely from appearance.
Changed check: Give a ligand-exchange example and ask whether colour change necessarily proves redox.
Case 18 — An extra reducing species can consume titrant
Question: An unknown solution contains Fe²⁺ and an unspecified other reagent that may also be oxidised.
Reasoning: The observed manganate titre could include reaction with more than Fe²⁺, so its volume alone may not uniquely determine iron concentration. A correct calculation needs the stated sample composition or evidence that the reagent is selective under the conditions. Inventing pure Fe²⁺ without justification would overclaim.
Changed check: Ask what additional analytical information or control would be needed before reporting a unique Fe²⁺ concentration.
Case 19 — A limiting reagent changes the theoretical yield
Question: The overall reaction equation is correct, but one redox reagent is present in insufficient amount.
Reasoning: Stoichiometry determines how far the reaction can proceed before that reagent is consumed. If the question asks for remaining Fe²⁺ or oxidant after mixing rather than an endpoint titre, use a mole ledger and identify the limiting species. The measured quantities might then describe an excess-reagent problem, not a standard equivalence calculation.
Changed check: Change the oxidant amount and ask which species remains after all possible reaction occurs.
Case 20 — The oxidising agent gains electrons
Question: A multiple-choice question asks whether manganate(VII) or Fe²⁺ is the oxidising agent.
Reasoning: In the standard acidified reaction, MnO4− gains electrons to become Mn²⁺ and is the oxidising agent. Fe²⁺ loses an electron to form Fe³⁺ and therefore functions as the reducing agent. The name oxidising agent describes what it causes to another reactant while undergoing reduction itself, a frequent source of reversed labels.
Changed check: Write oxidation-state changes explicitly for both species before selecting the answer option.
Case 21 — Check total charge on the full reaction
Question: A written redox equation appears to balance atoms but not ionic charge.
Reasoning: For MnO4−+5Fe²⁺+8H+→Mn²⁺+5Fe³⁺+4H2O, the total left charge is +17 and the right is also +17. Charge conservation is as essential as atom conservation. If an equation fails one of those checks, do not use its coefficients for quantitative work without correction.
Changed check: Construct the analogous dichromate equation and perform the total-charge check.
Case 22 — The best correction names the first mistake
Question: A student reaches a concentration exactly five times too small.
Reasoning: That pattern may indicate failure to multiply manganate moles by five to obtain Fe²⁺ moles, but the full working must be inspected rather than guessed. If the electron ratio is indeed missing, correct the half-equation and then set a different numerical example. The aim is to repair the chemistry, not simply multiply the final answer by five without understanding why.
Changed check: Have the student solve an oxalate example, where the correct factor is different, to verify transfer.
Case 23 — A final JC2 theory-to-practical comparison
Question: A learner knows all redox equations but struggles to describe the practical measurements that produce a titre.
Reasoning: A complete educational model distinguishes the measured delivered volume, the titrant’s known concentration, the unknown analyte amount and the electron-transfer ratio. Real practical competence also includes safe manipulation, observation and accurate recording under supervision; desk-based Chemistry teaching can support the reasoning but cannot replace authorised laboratory practice.
Changed check: Ask the student to interpret a fictional results table and separately identify which skills would need hands-on school training.
Case 24 — An unseen mixed problem
Question: A new exam question provides a Fe²⁺ aliquot, manganate concentration, two burette readings, a dilution history and asks for original stock concentration.
Reasoning: The dependable order is delivered titre from readings, titrant moles, Fe²⁺ moles via the five-to-one ratio, aliquot concentration, then reversal of the stated dilution factor to infer stock concentration. Each relationship represents a different stage, and none can be replaced by an arbitrary shortcut. The answer should include suitable units and numerical precision.
Changed check: Change one volume and one concentration and repeat the full method without copying the original numbers.
Frequently asked redox titration questions
How many Fe²⁺ ions react with one MnO₄⁻ under acidic conditions? Five. Why? Manganate accepts five electrons while each Fe²⁺ loses one. How many oxalate ions correspond to two MnO₄⁻? Five in the balanced acidic reaction. Does dichromate use the same Fe²⁺ ratio? No; one dichromate ion requires six electrons.
What is a titre? The delivered volume, calculated from burette readings. Do concordant readings prove accuracy? Not necessarily. Can a titre alone determine an unknown concentration? Only with appropriate titrant concentration, balanced chemistry and sample information. Are home redox experiments suitable? No; supervised laboratory facilities are required.
Further practice and JC2 enquiries
Continue with JC1 redox titration foundations, H2 manganate and dichromate titrations, Paper 4 planning and evaluation and the official 2027 syllabus. These resources cover complementary science and examination skills.
For class enquiries, contact eduKate Singapore or ask about current JC2 small-group arrangements. Confirm availability, fees, lesson location, timetable and practical training claims. Calculations here are original teaching examples, not official SEAB questions.