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A redox titre relates an oxidant to a reductant. For iron(II) and manganate(VII) in suitable acid, Fe²⁺ → Fe³⁺ + e⁻ and MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O combine to give a 5:1 Fe²⁺:MnO₄⁻ ratio. The acid is part of the balanced reaction, not an optional solvent label.
Video by Pdst Chemistry. Use the demonstration alongside this lesson’s specification reference and written explanations.
Viewing focus: Explain suitable conditions and end points in redox titrations
After watching, attempt these questions before opening the answers. You can also use the written lesson below.
After watching 1: Why can the iron determination use manganate as its own indicator?
Manganate is intensely coloured and is consumed while iron(II) remains. A faint persistent excess colour appears just after the iron(II) has reacted.
After watching 2: Explain the 5:1 iron(II):manganate ratio without quoting it from memory.
Each Fe²⁺ loses one electron, while each MnO₄⁻ gains five electrons in acid to become Mn²⁺. Five iron ions therefore supply the five electrons.
What you'll be able to do
A redox titre relates an oxidant to a reductant. For iron(II) and manganate(VII) in suitable acid, Fe²⁺ → Fe³⁺ + e⁻ and MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O combine to give a 5:1 Fe²⁺:MnO₄⁻ ratio. The acid is part of the balanced reaction, not an optional solvent label.
Use the approved sulfuric-acid conditions. Hydrochloric acid can supply chloride that is oxidised by manganate(VII), while nitric acid can oxidise iron(II); either can change the analyte or consume reagent. Insufficient suitable acid can produce a different manganese product. A correct numerical ratio cannot rescue a reaction conducted under the wrong chemical conditions.
Manganate(VII) supplies its own intense colour. When it is added to iron(II), each early portion is reduced; a faint persistent pink colour shows a small excess after the iron(II) is consumed. Use consistent swirling and dropwise addition near that point. A dark purple excess is an overshoot, not a more certain accurate result.
An iodine–thiosulfate titration instead uses I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. Add starch near the pale-iodine end point under the prescribed method, then use disappearance of the blue starch–iodine colour. Prolonged complex formation at high iodine concentration can make the end point sluggish. The amount ratio is one iodine to two thiosulfate, which differs from the iron–manganate ratio.
Record raw burette readings, exclude the rough run from the accurate mean and apply the stated concordance criterion. Use the actual aliquot and any dilution factor when calculating the original concentration. If the sample has to be dissolved or transferred, quantify that preparation rather than assuming the flask contains the original amount automatically.
Iron(II) can oxidise on standing in air; iodine can be lost or react under unsuitable conditions. Those changes can bias an otherwise precise titre. For manganate added to a fixed iron aliquot, an overshot titre overestimates the amount of iron inferred from that volume. A fresh sample or controlled preparation addresses chemical change; simply repeating a stale sample does not restore its original composition.
Attempt these problems on paper before opening the worked reasoning. Explain your steps as well as giving a final answer.
Common mistakes to avoid
Key takeaways
Question 1: A 10.00 cm³ iodine aliquot requires 18.60 cm³ of 0.01000 mol dm⁻³ thiosulfate. Find the iodine concentration.
0.009300 mol dm⁻³.
Check the assessed scope and practical or mathematical requirements.
An external explanation or interactive activity to use alongside the lesson. It may include extension material.
Close the notes and answer from memory. Revisit any explanation you cannot yet reconstruct.
Test yourself
Answer a few questions, then review anything you missed.
Start with the quiz above to check the key ideas. You can also choose another topic from the practice library.