Questions based on carrying out titration and interpreting the results are common both at AS and A level. Before we get onto the calculations, you need to have an understanding of what a titration is, how you make up the solutions, the method and what concordant results look like.
I’m not going to reinvent the wheel here, so I have put together a list of links to excellent YouTube videos that show you the practical techniques. I’m hoping that you will have had a number of opportunities to make up standard solutions and carry out titrations in the lab, so this should all be familiar.
| Making a standard solution |
| Carrying out a titration |
| Serial dilutions of a standard solution |
| Using a measuring cylinder |
| Using a volumetric pipette |
| Using a balance |
You need to be able to describe how to carry out a titration in detail, so make sure this is in your notes. You could use a method from a practical sheet, some text books have a section on practical skills, or you could use the mark scheme from a relevant exam question.
You also need to be able to explain how errors in carrying out a titration or making up a standard solution would affect the accuracy of calculations based on results from said titration.
Practice questions
1. Students titrating a 20.0 cm3 sample of an iodine solution against 0.00010 mol dm-3 sodium thiosulphate obtained an average titre of 5.30 cm3. It was suggested that the titre was too small and would lead to an unacceptably high uncertainty.
(a) Calculate the percentage error based on the students’ titre.
(b) Suggest how the experiment could be modified to improve the accuracy of the result.
2. A student makes up a standard solution of copper (II) sulphate by dissolving a weighed mass of solid in de-ionised water in a beaker. He transfers the solution into a volumetric flask, using a funnel, but he fails to wash out either the beaker or the funnel and transfer the washings into the flask before making up to the mark.
State and explain how the calculated concentration of the copper sulphate solution will be affected by this mistake.
3. A student plans to find the concentration of a solution of nitric acid by titrating it against a standard solution of potassium carbonate.
(a) Write a balanced equation, with state symbols for this reaction.
While making up the potassium carbonate solution, the student spilt some of the potassium carbonate on the bench as she tried to transfer it to a beaker from a weighing boat. Not all of the solid that had been weighed out was dissolved to make up the solution.
(b) Describe and explain how the concentration of the nitric acid, calculated from the results of the titration using this solution, will compare with the actual concentration of the acid.
(c) What piece of equipment would be used to be sure that when making up the standard solution exactly the right amount of water is added to the volumetric flask?
(d) Explain why the stoppered volumetric flask is inverted several times before the solution is used in the titration.
(e) Explain why the burette should be washed out with the potassium carbonate solution prior to use, and not with water.
4. The amount of iron (II) ions (Fe2+) in a compound can be determined through a redox titration in which the Fe2+ ions are oxidised by manganate (VII) ions in acidic solution, as shown below:
5Fe2+(aq) + MnO4–(aq) + 8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l)
The iron compound to be analysed was ground up, a certain mass weighed out, transferred to a beaker and sulphuric acid added to dissolve the compound. This was then carefully poured into a volumetric flask and the washings added. However, the student then overshot the graduation line when filling up the flask with sulphuric acid. He used a dropping pipette to remove the excess solution from the flask before stoppering and inverting.
(a) Describe and explain how the mass of iron in the compound calculated from the results of the titration might differ from the actual mass of iron in the compound.
(b) Suggest why the iron (II) compound is made up in acidic solution.
(c) The conical flask is rinsed with distilled water, but not dried, between titrations. Explain whether this will have any effect on the titres.
Answers:
1 (a) (0.1 / 5.30) x 100 = 1.89%
(b) Further dilute the concentration of the sodium thiosulphate solution or use a larger volume of iodine solution in each titration.
2 By failing to transfer washings to the volumetric flask not all of the copper sulphate mass that was weighed out is present in the standard solution – the actual concentration will be lower than the calculated concentration.
3 (a) K2CO3(aq) + 2HNO3(aq) ⇾ 2KNO3(aq) + H2O(l) + CO2(g)
(b) The potassium carbonate standard solution has a lower concentration than calculated as so it will take a smaller volume of nitric acid (a lower titre) to neutralise it. This means that the calculated concentration of the nitric acid will be higher than the actual concentration.
(c) Pipette or dropper
(d) To ensure than the solution is thoroughly mixed and therefore homogenous
(e) The burette is washed through with potassium carbonate solution before use to ensure it is not contaminated – water cannot be used as it would slightly dilute the potassium carbonate solution when the burette was filled.
4 (a) Some of the Fe2+ ions would have been removed so this standard solution is less concentrated than calculated. This means a lower volume of MnO4– will be needed to oxidise the Fe2+ and the mass of Fe present will be calculated to be lower than it is actually.
(b) It is made upon acidic solution because the equation for the oxidation shows the reaction requires H+ ions.
(c) This will have no effect on the titre because it doesn’t change the number of moles of Fe2+ pipetted into the conical flask for each titration.