How to Perform an Iron(II) Permanganate Redox Titration
A redox titration of iron(II) with potassium permanganate is a reliable way to determine the concentration of Fe²⁺ in a solution. In acidic conditions, permanganate ions oxidise iron(II) to iron(III), while manganese is reduced from oxidation state +7 to Mn²⁺. The reaction produces a strong visual endpoint, so no separate indicator is needed.
This experiment is common in analytical chemistry classes and teaching laboratories across Australia, including universities in Sydney, Melbourne, Brisbane and Perth. It is also a useful exercise for learning burette technique, stoichiometric calculations, uncertainty and safe handling of corrosive reagents. Broader educational resources for chemistry learners are available through NepaChemistry resources.
Aim And Reaction Chemistry
The purpose is to calculate the concentration of an iron(II) solution by measuring the volume of standard potassium permanganate solution required to reach the endpoint. The permanganate solution is placed in the burette, while a measured aliquot of Fe²⁺ is transferred to a conical flask.
The balanced ionic equation is:
[ \ce{MnO4^- + 5Fe^{2+} + 8H+ -> Mn^{2+} + 5Fe^{3+} + 4H2O} ]
One mole of permanganate reacts with five moles of iron(II). Sulfuric acid is preferred because sulfate ions do not undergo a significant redox reaction under these conditions. Hydrochloric acid can introduce chloride oxidation, while nitric acid is itself an oxidising agent.
Equipment And Reagents
You will need a burette, stand and clamp, volumetric pipette, pipette filler, conical flask, funnel, wash bottle, beaker and white tile. Use class A volumetric glassware where possible, particularly when the result will be assessed quantitatively. A white background makes the pale pink endpoint easier to observe.
The main reagents are the iron(II) sample, acidified potassium permanganate and dilute sulfuric acid. Potassium permanganate is a strong oxidiser and can stain skin, benches and clothing. Wear a laboratory coat, splash-resistant safety glasses and suitable gloves, and follow the current Safety Data Sheet and workplace health and safety procedures used in Australian laboratories.
Preparing The Sample
Rinse the burette with small portions of permanganate solution, then fill it slightly above the zero mark. Remove the funnel before titration, clear any air bubble from the jet and record the initial reading at eye level. Read the bottom of the meniscus and record readings to the appropriate precision for the burette.
Use a pipette to transfer a fixed volume, such as 25.00 mL, of the iron(II) solution into a clean conical flask. Add a measured volume of dilute sulfuric acid. Iron(II) solutions oxidise readily in contact with air, and this can happen more quickly during hot Australian summer conditions, especially in an inadequately sealed container. Prepare the sample promptly and keep it closed between measurements.
Standardising The Permanganate
Potassium permanganate is not usually treated as a primary standard because solid samples may contain impurities and solutions can slowly change concentration during storage. A laboratory may standardise it against a suitable primary standard, commonly sodium oxalate, before analysing the iron(II) sample.
The standardisation should be performed using the same acid concentration, temperature range and endpoint method as the main titration. Store permanganate in a clean amber bottle and protect it from strong light. In Australian university laboratories, prepared solutions are generally labelled with concentration, date, hazard information and preparer initials in line with local laboratory practice.
Performing The Titration
Place the acidified iron(II) sample beneath the burette and add permanganate while swirling continuously. At the beginning, the purple colour disappears quickly because each drop reacts with excess Fe²⁺. As the endpoint approaches, add the titrant dropwise, washing the inside wall of the flask with deionised water when necessary.
The endpoint is the first faint pink colour that persists for about 30 seconds. It should be a very pale tint, not a deep purple colour. Record the final burette reading and calculate the titre by subtracting the initial reading. Repeat the titration until at least two concordant titres agree within the laboratory’s stated tolerance, often about 0.10 mL.
Calculations And Quality Checks
First calculate the amount of permanganate used:
[ n(\ce{MnO4^-}) = c(\ce{MnO4^-}) \times V(\ce{MnO4^-}) ]
Use litres for volume in this equation. From the reaction ratio, the amount of iron(II) is:
[ n(\ce{Fe^{2+}}) = 5n(\ce{MnO4^-}) ]
The iron(II) concentration is then:
[ c(\ce{Fe^{2+}})=\frac{n(\ce{Fe^{2+}})}{V(\ce{Fe^{2+}\ sample})} ]
For example, if 25.00 mL of sample requires 18.60 mL of 0.02000 mol L⁻¹ permanganate, the sample contains 0.001860 mol Fe²⁺, giving a concentration of 0.07440 mol L⁻¹. Keep extra digits during the calculation and round the final result according to the uncertainty of the measurements.
Results are weakened by overshooting the endpoint, reading the burette at an angle, using insufficient acid or allowing the iron(II) solution to oxidise before titration. Rinsing the conical flask with deionised water does not affect the result, but rinsing the pipette with water would dilute the sample.
Choosing Suitable Conditions
Different approaches can be appropriate depending on whether the goal is a teaching demonstration, a formal analysis or a rapid screening measurement. The following comparison highlights the practical trade-offs.
| Feature | Permanganate Titration | Dichromate Titration | Instrumental Iron Analysis |
|---|---|---|---|
| Endpoint | Self-indicating pale pink | Requires an indicator | Instrument signal |
| Main reagent concern | Strong oxidiser and staining | Toxic chromium(VI) waste | Instrument and calibration needs |
| Equipment | Basic volumetric glassware | Basic glassware plus indicator | Spectrophotometer or atomic instrument |
| Best use | Teaching and routine Fe²⁺ analysis | Controlled analytical laboratories | Low concentrations or complex samples |
| Waste handling | Oxidising manganese-containing waste | More demanding hazardous waste control | Method-dependent chemical waste |
For a school or undergraduate laboratory in Australia, permanganate is often selected because the colour change is easy to see and the apparatus is inexpensive. However, laboratories must still collect waste correctly rather than pouring concentrated reagent into a sink. Reagent availability varies between local scientific suppliers, and household products bought through the Australian retail market are not suitable substitutes for analytical-grade chemicals.
Recording A Defensible Result
A complete report should include the standardised permanganate concentration, sample volume, individual titres, concordant titres, mean titre, balanced equation and final Fe²⁺ concentration. State the acid used and describe the endpoint clearly. If the work is part of a regulated analysis, use the laboratory’s validated method and quality-control requirements; NATA-accredited facilities may require additional traceability and documentation.
The most dependable workflow is simple: acidify a fresh iron(II) aliquot with sulfuric acid, titrate slowly with standardised permanganate until a faint persistent pink colour remains, repeat for concordant results, and apply the 1:5 stoichiometric ratio carefully. Linked burette readings, controlled acidity and prompt handling are the practical foundations of a sound result.