11 – Standardising a Potassium Permanganate Solution (LC 2027) (Leaving Cert Chemistry): Revision Notes

11 – Standardising a Potassium Permanganate Solution

Introduction to standardisation

Standardisation is the process of determining the exact concentration of a solution by titrating it against a substance of known concentration called a standard solution. In this experiment, you'll prepare a standard solution of ammonium iron(II) sulphate and use it to find the precise concentration of a potassium permanganate solution through titration.

Preparing the standard solution

The first crucial step involves creating an accurate standard solution of ammonium iron(II) sulphate. This compound serves as an excellent primary standard because it's stable, doesn't absorb moisture from the air, and has a high molar mass that minimises weighing errors.

You begin by accurately weighing the required mass of ammonium iron(II) sulphate crystals using a clock glass and laboratory balance. The precision of this weighing is vital since any error will carry through to your final results.

Next, you'll dissolve the crystals in dilute sulfuric acid rather than pure water. This acid serves a critical purpose: it prevents the iron(II) ions from being oxidised to iron(III) ions by oxygen in the air. Without this protection, your Fe²⁺ ions would convert to Fe³⁺ ions, making them unavailable for the redox reaction with permanganate.

The dissolved solution is then carefully transferred to a volumetric flask using a funnel, ensuring all washings from the beaker are included. After diluting to the graduation mark with deionised water and inverting the flask about 20 times, you have a standard solution of known, precise concentration.

Setting up the titration

Proper preparation of your glassware is essential for accurate results. Begin by washing all equipment - pipette, burette, and conical flask - with deionised water. The burette requires special attention: rinse it with some of the potassium permanganate solution to ensure the concentration isn't diluted by residual water.

Fill the burette with permanganate solution using a funnel, then remove the funnel before taking readings. Adjust the solution level so the top of the meniscus aligns with the zero mark, remembering to read at eye level for accuracy. Crucially, ensure the space below the tap is filled with solution to avoid air bubbles affecting your volume measurements.

Carrying out the titration procedure

Using a pipette filler, transfer a precise volume of your standard ammonium iron(II) sulphate solution into the conical flask. Add excess dilute sulfuric acid to this solution - the acid must be in excess to provide sufficient hydrogen ions for the redox reaction.

The chemical reaction requires hydrogen ions according to this half-equation:

$\text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}$

Begin titrating by adding permanganate solution from the burette while constantly swirling the conical flask. You'll notice an interesting phenomenon called autocatalysis: the first few drops of permanganate are decolourised slowly, but as more manganese(II) ions form, they catalyse the reaction, causing subsequent drops to react much more rapidly.

Understanding the end point

Potassium permanganate is described as self-indicating, meaning it acts as its own indicator. The deep purple permanganate solution loses its colour when it reacts with iron(II) ions, but once all the iron(II) has been consumed, the next single drop of permanganate will impart a permanent faint pink colour to the solution.

This colour change marks your end point - the moment when equivalent amounts of permanganate and iron(II) have reacted. The permanganate essentially indicates its own excess, eliminating the need for additional indicators.

You should perform one rough titration followed by two accurate titrations. The accurate results should agree within 0.1 cm³, demonstrating the precision of your technique. This concordance gives confidence in your results.

Chemical reactions and equations

The overall reaction involves the reduction of permanganate ions by iron(II) ions in acidic solution. The permanganate acts as an oxidising agent, accepting electrons from the iron(II) ions:

Reduction half-equation: $\text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}$

Oxidation half-equation: $\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + \text{e}^-$

Overall balanced equation: $\text{MnO}_4^- + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} + 5\text{Fe}^{3+}$

The stoicheiometry shows that one mole of permanganate ions reacts with five moles of iron(II) ions, which is crucial for your calculations.

Calculations and results

Using your average concordant titre volume, you can calculate the exact molar concentration of the permanganate solution. The calculation involves using the mole ratio from the balanced equation (1:5 ratio of MnO₄⁻ to Fe²⁺) and the known concentration and volume of your standard iron(II) solution.

These calculations allow you to determine the precise molarity of your permanganate solution, which can then be used for further quantitative analysis experiments.

Important safety considerations

Common issues and troubleshooting

Summary