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Iron tablets may be used in the treatment of anemia - Leaving Cert Chemistry - Question 1 - 2003
Question 1
Iron tablets may be used in the treatment of anemia. To analyse the iron(II) content of commercially available iron tablets a student used four tablets, each of mass... show full transcript
Worked Solution & Example Answer:Iron tablets may be used in the treatment of anemia - Leaving Cert Chemistry - Question 1 - 2003
Step 1
a) Why was it important to use dilute sulfuric acid as well as deionised water in making up the solution from the tablets?
Answer
It was important to use dilute sulfuric acid to prevent the oxidation of iron(II) ions to iron(III) ions by air. The acid helps to maintain the iron in its reduced form, ensuring accurate titration results.
Step 2
b) Describe in detail the procedure for making up the 250 cm³ solution from the tablets.
Answer
To prepare the solution:
- Crush the iron tablets using a mortar and pestle until a fine powder is achieved.
- Transfer the crushed tablets into a beaker.
- Rinse the mortar with deionised water and add this to the beaker to ensure all solids are transferred.
- Measure 250 cm³ of deionised water in a graduated cylinder.
- Add about 15 cm³ of dilute sulfuric acid to the beaker containing the powdered tablets.
- Stir the mixture to dissolve the tablets fully.
- Transfer the solution into a volumetric flask carefully, using a funnel if necessary.
- Rinse the beaker and funnel with deionised water to ensure all solution is transferred.
- Finally, make up the volumetric flask to the mark with more deionised water, ensuring accurate concentration.
Step 3
c) Why was more dilute sulfuric acid added before the titrations were commenced?
Answer
Additional dilute sulfuric acid was added to ensure that the iron(II) ions remained in solution and did not oxidize to iron(III) ions during the titration process. This ensures that the titration reaction accurately reflects the amount of iron(II) present.
Step 4
d) How was the end-point detected?
Answer
The end-point of the titration was detected by observing a color change. Initially, the solution is pale yellow due to the presence of iron(II) ions. When all iron(II) has reacted and the first excess of potassium manganate(VII) is added, the solution turns a permanent pink (purple) color due to the formation of manganese(II) ions.
Step 5
e) (i) the concentration of iron(II) solution in moles per litre
Answer
First, calculate the number of moles of KMnO₄ used:
Number of moles = Concentration × Volume = 0.010 mol/L × 0.0139 L = 0.000139 mol.
From the balanced equation: 1 mole of MnO₄⁻ reacts with 5 moles of Fe²⁺. Thus, Number of moles of Fe²⁺ = 5 × 0.000139 mol = 0.000695 mol in 25 cm³.
To find the concentration in 250 cm³:
Concentration = Number of moles / Volume(L) = 0.000695 mol / 0.025 L = 0.0278 mol/L.
Step 6
e) (ii) the mass of iron(II) in one tablet
Answer
From the concentration calculated: Mass of iron(II) = Number of moles × Molar mass (Fe = 55.85 g/mol).
In 250 cm³, total moles of Fe²⁺ = 0.0278 mol/L × 0.250 L = 0.00695 mol.
Thus, mass of Fe²⁺ = 0.00695 mol × 55.85 g/mol = 0.388 g.
Since this is from 4 tablets, mass of iron(II) per tablet = 0.388 g / 4 = 0.097 g.
Step 7