Understand the range of different titration procedures that 16–18 students may need to be familiar with

Preparing a standard solution

It is important to be able to know how to create a standard solution before you can even get started with a titration. A standard solution needs to be accurately prepared to give a known concentration of a dissolved element or compound. This video outlines the procedure for preparing a standard solution and how to calculate the final concentration. Note the use of the glass rod to aid with transferring the solution from the beaker to the volumetric flask.

Questions you can ask your students:

  • What method is used to measure the mass of the sodium carbonate?
  • Why is weighing by difference more accurate than other alternatives?
  • Why is it important to rinse the glass rod and beaker and transfer the washings to the volumetric flask?

Students can also be asked to set questions for each other to determine the concentrations of various standard solutions. One of the easy mistakes in these multi-step calculations is using volumes expressed in cm3 and not dm3.

Acid-base neutralisation

One of the most common titrations is an acid-base neutralisation, where the concentration of one of the reactants is determined via titration. In the experiment below, methyl orange is used to identify the endpoint of the titration, but some other titrations will use phenolphthalein. The video below outlines the titration procedure. Note that as the titration nears its endpoint, distilled water is used to wash down the burette tip and the sides of the conical flask to ensure that all reactants are mixed together.

Questions you can ask your students:

  • Why don’t we use Universal Indicator in acid-base titrations?
  • Why is the conical flask placed on a white tile?
  • Why is it useful to carry out a rough or ‘overshoot’ titration?

The associated calculations for an acid-base titration (in this case, hydrochloric acid and sodium hydroxide) are shown in the following video. Ask your students to write a balanced equation of the equation before attempting any calculations. In titrations, it can be easy to lose track of which reactant was in the burette and which was in the conical. When writing out calculations, encourage your students to include helpful details. For example, Moles of Na2CO3 titrated into conical flask may be more helpful than simply writing Moles of Na2CO3. In addition, students need to aware that some acids – such as sulfuric acid – are diprotic, releasing two H+ ions in solution.

Suggested questions:

  • What do we mean when we say, ‘concordant titre’?
  • In molar calculations, how do we determine the number of significant figures to use when recording our result?

The following video demonstrates a neutralisation titration using phenolphthalein as the indicator. This video demonstrates a better hand technique for dispensing from the burette with the fingers approaching the tap from behind rather than the side. If teaching this practical remotely, highlight to students that the pipette is allowed to drain under gravity and is calibrated so that a small volume of solution will remain in the tip. This is something that is likely to go overlooked without live practical experience.

Redox Titration

Acid-base neutralisation titrations are one of the more straightforward titrations, but they are an excellent example for helping students become familiar with the technique. Another common procedure is a redox titration, where one reactant is oxidised while the other is reduced. In this video from SpaceyScience potassium permanganate is used to determine the mass of iron(II) in a salt (or an iron tablet in other similar experiments). The iron ion oxidises while the manganate ions are reduced as shown in the equation below:

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

MnO4- ions in solution are purple while Mn2+ ions are colourless. The endpoint is determined when the iron solution in the conical flask turns purple, indicating that all the iron(II) has been oxidised and the excess MnO4- is no longer reduced, retaining its purple colour.

Determining the mass of iron in the salt is not simple, as it requires multi-step calculations which incorporate the 1:5 ratio for Mn:Fe ions in the balanced equation. The titration screen experiment is a helpful resource to support students to work through the calculations.

Another redox titration involves titrating sodium thiosulfate into an unknown quantity of iodine using starch as an indicator. This short video from KEGS Chemistry demonstrates the colour change and how the titration could be used to determine the concentration of Fe(III) ions in a solution using a back titration.

Back Titration and Double Titration

A number of specifications may also require students to have an understanding of ‘back’ and ‘double’ titrations, with the latter technique requiring the use of two indicators. These two videos from SpaceyScience provide further examples of possible titrations students may encounter. You can find an explanation of double titration in this pdf.

Also check out…

  • Titration screen experiment – an online experiment to familiarise students with the titration procedure and a range of applications, from a simple acid-base neutralisation to a more extended redox titration.

  • Quantitative chemistry Starter for 10 – use these sets of titration calculations as a diagnostic tool or as a plenary to consolidate learning.

  • A thermometric titration – it’s possible to use temperature to determine the endpoint of a reaction rather than an indicator such as methyl orange or phenolphthalein.

  • A conductimetric titration using acids and alkalis – the changes in ion concentrations during a neutralisation reaction can be observed using a simple circuit and carbon electrodes.

  • Sulfur dioxide in wine – this redox titration uses iodine and starch as an indicator and determines the concentration on a sulfur dioxide in a wine sample. You can also watch a video here.

  • Moles and titrations: scary stuff? – this article tackles some of the challenges and misconceptions that students encounter in titration calculations.