You have likely performed a titration before. Perhaps you used it to find out how much of an alkali you need to neutralise a certain amount of acid. Redox titrations help us find the exact amount of an oxidising agent needed to react with a reducing agent.

Redox titrations with transition metals are exciting because of their colourful variable oxidation states. Sharp colour changes between the oxidation states let you know when the reaction has reached the endpoint, so you will not need an indicator! Let us use potassium manganate(VII) as an example to find out how this works!

• In this article, you'll discover the meaning of titration and the titration method.
• You will perform the redox titration of manganate(VII) with iron.
• We will also examine the redox titration of manganate(VII) with ethanedioate ions.
• You will learn how to do calculations based on these two titrations.

Titration meaning

Titration is a way of analysing chemicals to find an unknown concentration by using a substance with a known concentration.

Essentially, we slowly add a standard solution of a titrant from a burette to the analyte in the conical flask. We can use a colour indicator in order to know that the reaction has reached its endpoint.

Performing a titration will help you understand how it works. Let us next examine the steps involved in a titration.

Titration method

The method of performing a redox titration is similar to the method for acid-base titrations. You may read about it in pH Curves and Titrations.

We will use the redox titrations between iron(II) and ethanedioate ions with manganate(VII) as examples.

Redox titration of iron(II) with manganate(VII)

People who suffer from anaemia - low iron concentration in the blood - may be prescribed iron tablets by their doctor or pharmacist. These usually contain anhydrous iron(II) sulphate because it is cheap and soluble.

You can estimate the amount of iron(II) sulphate in each tablet by titrating it against a standard solution of potassium manganate(VII). You will have to dissolve each tablet in diluted sulfuric acid first!

In a laboratory, you may carry out the following experiment with iron tablets from the pharmacist.

1. Weigh out 8 iron tablets. Dissolve them in a beaker of about 100 cm³ of 2 M sulfuric acid.
2. Some tablets have an outer coating which may not dissolve. You may need to filter the solution.
3. Pour the filtrate into a 250cm³ volumetric flask. Ensure you include the washings from the conical flask and filter paper. Use distilled water to fill up to the mark.
4. Pipette 25 cm³ of the solution into a conical flask. Add 25 cm³ of dilute sulfuric acid.
5. Titrate against 0.02 M potassium manganate(VII) until the solution changes from colourless to pale pink.
6. Repeat the experiment until you get a concordance of ±0.10 cm³.

A diagram of the equipment you will need is shown below.

Diagram of titration equipment and method, Olive [Odagbu] Vaia Originals

In this reaction, Fe²⁺ gets oxidised to Fe³⁺ while Mn⁷⁺ gets reduced to Mn²⁺. You write the half equations for the process as follows:

Redox reactions between manganate(VII) and iron(II), Vaia Originals

You must use diluted sulphuric acid because potassium permanganate works best as an oxidiser in acidic conditions. Remember, transition metal ions require strongly acidic conditions when going from a higher to a lower oxidation state. However, you cannot use just any acid!

• Manganate(VII) oxidises chloride ions in hydrochloric acid to chlorine.
• Weak acids like ethanoic acids do not provide enough H⁺ ions.
• Using a concentrated sulphuric acid or nitric acid may oxidise the analyte.

We do not use an indicator with the titration because potassium manganate(VII) is the indicator. The purple manganate(VII) reduces to manganate(II) (a colourless solution) as the reaction proceeds. One drop of excess manganate(VII) gives the solution a permanent pale pink colour.

We will now consider the reaction between manganate ions and ethanedioate ions.

Titrations with ethanedioate ions

The reaction between manganate and ethanedioate ions (C₂O₄^2-) is intriguing because it is autocatalytic. Chemists use ethanedioic acid (also called oxalic acid) to standardise or determine the strength of permanganate solution.

The redox reaction between manganate(VII) and ethanedioate ions takes place as follows:

MnO₄^- is reduced to Mn²⁺ and C₂O₄^2- is oxidised to CO₂.

Redox reaction between manganate (VII) and ethanedioate ions, Vaia Originals

Here are the steps to perform the titration:

1. Rinse and fill a clean burette with the potassium permanganate solution.
2. Attach the burette to the burette stand and place the white tile below the conical flask.
3. Pipette out 10cm³ of 0.1M ethanedioic acid solution into a clean conical flask.
4. Add 5cm³ of dilute sulfuric acid to the conical flask.
5. Heat the ethanedioic acid solution to 60ºC.
6. Note down the initial reading in the burette.
7. Titrate the hot ethanedioic acid solution against the potassium permanganate solution whilst continuously swirling the flask gently.
8. Stop when you observe a permanent pale pink colour solution.
9. Record the reading from the upper meniscus on the burette.
10. Repeat the experiment until you get a concordance of ∓0.10cm³.

We heat the ethanedioate solution to about 60-70ºC to speed up the reaction with potassium permanganate. Be careful not to heat the solution past 70ºC, as ethanedioate begins to decompose at 70ºC and above.

Permanganate's oxidising power works best in an acidic environment. So we use dilute sulfuric acid in this experiment. Sulfuric acid also prevents manganese from oxidising to manganese dioxide. We cannot carry out the titration in the presence of acids such as hydrochloric acid or nitric acid. Hydrochloric acid is an oxidising agent that reacts with manganate(VII) to form chlorine.

As with the previous titration, permanganate acts as a self indicator. Purple MnO₄^- ions reduce to colourless Mn²⁺ ions. One drop of excess MnO₄^- ions presents a pale pink colour.

Have you got that? Okay, let us do some calculations!

Titration calculations

After you have completed a titration, you will need to do some calculations to determine the concentration of the analyte. Let us try a few together!

Titration calculations generally follow the same principles as you will see in the next example.

It takes practice to get the hang of titration calculations. Try the examples in the exercises section to improve your skills!