Chemistry Life
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
w w w e tr .X
Appendix 2
m eP
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.
e ap .c rs om
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