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Titrations

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Titrations

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.

  1. A titrant is the name of the chemical of known concentration.
  2. We call the unknown substance the analyte.
  3. A standard solution is a solution whose exact concentration is known.

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.

Titrations, Diagram of method and equipment, StudySmarterDiagram of titration equipment and method, Olive [Odagbu] StudySmarter 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:

Titrations, Redox Managanate(VII) and Iron(II), StudySmarterRedox reactions between manganate(VII) and iron(II), StudySmarter 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 (C2O42-) is intriguing because it is autocatalytic. Chemists use ethanedioic acid (also called oxalic acid) to standardise or determine the strength of permanganate solution.

Read about autocatalysts in Catalysts.

Ethanedioic acid, also called oxalic acid, can be found in plants such as spinach and rhubarb. Oxalic acid salts contain the ethanedioate ion (C2O42-). We can learn the concentration of free oxalate ions in solution by titrating against potassium permanganate. This reaction is used to analyse the ethanedioate content of spinach leaves, for example.

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

MnO4- is reduced to Mn2+ and C2O42- is oxidised to CO2.

Titrations, Redox manganate (VII) and ethanedioate ions, StudySmarterRedox reaction between manganate (VII) and ethanedioate ions, StudySmarter 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!

24.55cm3 of 0.020M aqueous potassium manganate(VII) reacted with 25.0cm3 of acidified iron(II) sulfate solution. Find the concentration of Fe2+ ions in the solution.

Step 1: Write out the balanced equation

5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)

Step 2: Work out the number of moles of MnO4- ions added to the flask.

Moles of MnO4- =

We divide by 1000 to convert the volume from cm3 to dm3.

Moles of MnO4- =

Step 3: The equation tells you that 1 mole of MnO4- reacts with 5 moles of Fe2+.

5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)

Step 4: Multiply moles of MnO4- by 5.

0.000419 x 5 = 0.002455 moles of Fe2+

Step 5: Work out the concentration of Fe2+

moles of Fe2+ =

Rearrange so that concentration =

Concentration =

Concentration = 0.0982 mol dm-3

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

Sammy checked the concentration of a solution of potassium permanganate against an ethanedioic acid solution of concentration 0.04 mol dm-3.

He placed 25cm3 of the ethanedioic acid solution in a flask with excess dilute sulphuric acid. After warming the solution, he carried out a titration. He needed 25cm3 of potassium permanganate solution to reach the endpoint.

Calculate the actual concentration of the permanganate solution.

Step 1: Write the balanced equation for the reaction

2MnO4 + 16H+ + 5C2O42- → 2Mn2+ + 10CO2 + 8H2O

Step 2: Find the number of moles of ethanedioic acid

No. of moles of 5C2O42- =

Step 3: Find the number of moles of potassium manganate(VII).

The balanced equation tells us that we need as many moles of permanganate ions as ethanedioate ions.

2MnO4- (aq) + 16H+ (aq) + 5C2O42- (aq) ➔ 2Mn2+ (aq) + 10CO2 (aq) + 8H2O (l)

Step 4: Multiply the no. of moles of ethanedioate ions by

0.001 x in 25cm3 = 0.0004 moles of manganate (VIII)

Step 5: Find the concentration by rearranging the formula

Rearrange the formula so that

Concentration = 0.0004 x

Concentration = 0.016 mol dm-3

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

Titrations - Key takeaways

  • You can estimate the amount of iron(II) sulphate in each tablet by titrating it against a standard solution of potassium manganate(VII).
  • In the redox reaction between manganate(VII) and iron(II), Fe2+ is oxidised to Fe3+ while Mn7+ is reduced to Mn2+.
  • Redox titrations of iron(II) and ethanedioate ions with manganate(VII) must be acidified with dilute sulfuric acid.
  • Redox titrations of iron(II) and ethanedioate ions with manganate(VII) cannot be acidified with hydrochloric acid because it is oxidised to chlorine. Ethanoic acid is too weak and does not provide enough H+ ions. Nitric acid, as well as concentrated sulphuric acid, oxidise the analyte.
  • 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.
  • In the redox titration between permanganate and ethanedioic acid, we heat the ethanedioic acid solution to about 60-70ºC to speed up the reaction with potassium permanganate.
  • Adding sulphuric acid to the analyte in titrations with permanganate prevents manganese from oxidising to manganese dioxide.
  • The redox reaction between manganate(VII) and ethanedioate ions takes place as follows: 2MnO4 (aq) + 16H+ (aq) + 5C2O42- (aq) ➔ 2Mn2+ (aq) + 10CO2 (aq) + 8H2O (l).

Frequently Asked Questions about Titrations

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

Universal indicator gives a different colour for different pH ranges. That makes it hard to titrate to a specific pH value. On the other hand, specialised indicators like phenolphthalein change from colourless to deep red at pH above 9.0.

  1. Using a clean pipette, measure a set volume of a solution of unknown concentration into a clean conical flask.
  2. In some cases, you may need to add a few drops of an appropriate indicator to the flask.
  3. Place a white tile under the flask so you can easily see any colour changes.
  4. Rinse and fill the burette with the solution of known concentration.
  5. Record the starting volume in the burette.
  6. Slowly add the solution in the burette to the solution in the conical flask while gently swirling the flask.
  7. Stop the titration when you reach the endpoint. The endpoint is when one drop of excess solution from the burette changes the colour of the solution in the flask.
  8. Record the final volume from the burette.
  9. Subtract the initial burette reading from the final burette reading to obtain the titre.
  10. Repeat the titration until you get concordant titre values of ∓0.10 cm3.

Let us try an example together. 

23.9cm3 of 0.040 mol dm-3 of aqueous potassium permanganate reacted with 25cm3 of acidified iron(II) sulphate solution.  What was the concentration of Fe2+ ions in the solution?


First things first, write down the equation for the reaction. The redox process between manganate(VII) and iron(II) takes place as follows:


5Fe2+ (aq) + MnO4- (aq) + 8H+ (aq) ➔ 5Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)


Next, use the values provided to find the number of moles of MnO4- ions added to the flask.

Use the formula: no. of moles = concentration x volume / 1000


0.04 x 23.9/1000 = 0.000956 or 9.56x10-4 moles of MnO4-


Now we can figure out the number of moles of Fe2+ in the flask!

Use the reaction equation to find the proportion of the reaction between the titrant and analyte.

From the equation, we can see that 1 mole of manganate(VII) reacts with 5 moles of iron(II).

So multiply 9.56x10-4 by 5.


9.56x10-4 x 5 = 0.00478 or 4.78x10-3 of Fe2+ ions


Use the previous formula to calculate the concentration of Fe2+ ions.

0.00478 = Conc. x 25 / 1000


Conc. = 0.00478 x 1000 / 25


Conc. = 0.1912 mol dm-3


Congratulations, you have completed a titration calculation! Getting the hang of titration calculations takes practice. Have a go at the examples in the exercises section.

Final Titrations Quiz

Question

When trying to find the amount of iron(II) sulphate in an iron tablet, why might you have to filter the solution after you dissolve the tablets?

Show answer

Answer

You filter the iron(II) sulphate solution because some iron tablets have an insoluble outer coating.

Show question

Question

What colour change do you expect to see at the endpoint of a titration between potassium manganate(VII) and Fe(II)?

Show answer

Answer

Colourless to pale pink.

Show question

Question

 Write the half equations for the reaction between permanganate and iron(II).

Show answer

Answer

MnO4- + 8H+ + 5e- ➔ Mn2+ + 4H2O


5Fe2+ ➔ 5Fe3+ + 5e-


Show question

Question

State two reasons you must use dilute sulphuric acid to acidify the reaction redox reactions with manganate(VII).

Show answer

Answer

  • Concentrated sulphuric acid may oxidise the analyte
  • Sulphuric acid prevents manganese from oxidising to manganese dioxide
  • Manganate(VII) acts as a strong oxidiser in acidic conditions
  • To provide enough H+ ions

Show question

Question

Why do we not use an indicator in the redox titration between manganate(VII) and ethanedioic acid?

Show answer

Answer

Potassium permanganate acts as a self indicator.

Show question

Question

 Name the following ion: C2O42-.

Show answer

Answer

Ethanedioate or Oxalate ion


Show question

Question

Write the half equations for the reaction between permanganate and ethanedioate ions.

Show answer

Answer

MnO4- + 8H+ + 5e- ➔ Mn2+ + 4H2O


C2O42- ➔ 2CO2 + 2e-


Show question

Question

Why do we heat the ethanedioic acid solution before we titrate against permanganate?

Show answer

Answer

Heating speeds up the reaction between ethanedioate and manganate(VII).

Show question

Question

What must we heat ethanedioic acid solution to between 60 and 70ºC before titrating it against permanganate?

Show answer

Answer

Ethanedioate begins to decompose over 70ºC.

Show question

Question

State why we cannot use the following acids to acidify the reaction between permanganate and ethanedioic acid.

  • Hydrochloric acid
  • Nitric acid
  • Ethanoic acid


Show answer

Answer

  • Hydrochloric acid gets oxidised by manganate(VII) to chlorine
  • Nitric acid may oxidise the analyte.
  • Ethanoic acid will not provide enough H+ ions because it is a weak acid.

Show question

Question

What do we call the chemical of unknown concentration in a titration?

Show answer

Answer

The Analyte.

Show question

Question

A standard solution of potassium permanganate can be used to determine the concentration of free ethanedioate ions in solution. Outline the method for this experiment.

Show answer

Answer

  1. Rinse and fill a clean burette with the potassium permanganate solution.
  2. Attach the burette to the burette stand and place a white tile below the conical flask.
  3. Pipette out a set amount of ethanedioic acid solution into a clean conical flask.
  4. Add excess 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.


Show question

Question

To determine the amount of Fe2+ ions in iron tablets, Kelly dissolved them in hydrochloric acid and titrated them against potassium dichromate(VI). Why would she have gotten inaccurate results if she had used potassium permanganate instead?

Show answer

Answer

Hydrochloric acid oxidises to chlorine in the presence of potassium manganate(VII)

Show question

Question

Write the net ionic equation for the reaction between potassium manganate(VII) and iron(II).

Show answer

Answer

MnO4- + 8H+ + 5Fe2+ ➔ Mn2+ + 4H2O + 5Fe3+  


Show question

Question

Label the equipment and substances in the titration diagram below. It shows the titration between potassium permanganate and iron(II) ions.


Show answer

Answer

A: Burette

B: Potassium manganate(VII) or potassium permanganate

C: Clamp stand

D: Conical flask

E: Fe(II) sulphate solution

Show question

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