# Chemistry Life

Topics: Titration, Iron, Laboratory glassware Pages: 6 (1815 words) Published: December 10, 2012
Teaching AS Chemistry Practical Skills

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Appendix 2

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3. How much iron is there in an iron tablet? – Student Sheet In this practical you will have the opportunity to perform a quantitative analysis using the technique of titration. You are going to analyse an iron tablet to find out how much iron is actually present in it. Titrations involving potassium manganate(VII) may form part of your Practical Assessment. Intended lesson outcomes By the end of this practical you should be able to: • perform a titration involving potassium manganate(VII); • read a burette and use a pipette; • use a volumetric flask; • record your titration results appropriately in tables you have drawn yourself; • use and understand an ionic equation; • use the mole concept to perform calculations. Background information Iron performs a vital role in our bodies. It is present in red blood cells and forms part of the haemoglobin molecule, which combines with oxygen from the lungs. The oxygen is then transported all round the body. When young people are growing rapidly, the body may not have enough iron and this causes anaemia. This can be remedied by a course of ‘ferrous sulphate’ tablets, often known as iron tablets.

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The iron in iron tablets is in the form of hydrated iron(II) sulphate (sometimes called ferrous sulphate). As the name iron(II) suggests, the Fe2+ ion is present. To determine just how much Fe2+ is in each tablet, we can react the Fe2+ ions with manganate(VII) ions, which have a formula MnO4-. MnO4-(aq) + deep purple 8H+(aq) + 5Fe2+(aq) pale green → Mn2+(aq) + almost colourless 5Fe3+(aq) + brown 4H2O(l)

Although this ionic equation may appear complicated at this point in your course, you can see from the colours that the deep purple of the MnO4-(aq) will disappear when the reaction is complete. The end point is when the addition of one extra drop of potassium manganate(VII) solution turns the solution in the conical flask to a pale pink colour. Question 1 Explain why this volumetric analysis does not require an indicator.

Safety You must wear eye protection throughout this experiment.

1.0 mol dm-3 sulphuric acid is irritant. Wash all spillages with plenty of water. Always use a pipette filler, or other suitable safety device. Never be tempted to use your mouth to draw liquid into the pipette.

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© University of Cambridge International Examinations 2006

Teaching AS Chemistry Practical Skills Appendix 2 Procedure 1. Accurately weigh two iron tablets and record their mass. 2. Grind up the tablets with about 5 cm3 of 1.0 mol dm-3 sulphuric acid, using a pestle and mortar. Transfer this into a 100 cm3 volumetric flask. Use further small volumes of the dilute sulphuric acid to rinse the ground-up tablets into the flask, until no traces of the iron tablets are left in the mortar. Add more 1.0 mol dm-3 sulphuric acid to make the volume in the volumetric flask exactly 100.0 cm3. Stopper the flask and shake it thoroughly to mix the solution. Not all of the outer coating of the tablets will dissolve. This does not matter as it does not contain any Fe2+ ions. Use a pipette and pipette filler to withdraw 10.0 cm3 of the tablet solution and transfer it into a clean conical flask. Wash a 100 cm3 beaker with deionised water and then twice with small volumes of the 0.0050 mol dm3 potassium manganate(VII) solution you are going to use. Now half fill the beaker with the potassium manganate(VII) solution and use this to fill the burette. Remember that the burette does not have to be filled to the 0.00 cm3 mark. Make sure that you run some of the solution back into the beaker to ensure that the tip of the burette is full. Read the volume of potassium manganate (VII) solution in the burette. In this case the colour of the potassium manganate(VII) solution is so intense that you cannot see the bottom of the meniscus so you must use the top of the meniscus to measure the volume. Draw up a table in...