Introduction
Iron is one of the many minerals required by the human body. It is used in the manufacture of oxygen-carrying proteins, haemoglobin and myoglobin. A deficiency of iron in the body can leave a person feeling tired and listless, and can lead to a disorder called anemia. Many of the foods we eat contain small quantities of iron.
In this analysis, the iron present in an iron tablet (dietary supplement) or a sample of food is extracted to form a solution containing Fe3+ (ferric) ions. To make the presence of these ions in solution visible, thiocyanate ions (SCN−) are added. These react with the Fe3+ ions to form a blood-red coloured complex:
Fe3+(aq) + SCN−(aq) = [FeSCN]2+(aq)
By comparing the intensity of the colour of this solution with the colours of a series of standard solutions, with known Fe3+ concentrations, the concentration of iron in the tablet or food sample may be determined. This technique is called colorimetry.
Methods
Preparation of Fe3+ standard solutions
NB: It may take several days to dissolve the Fe3+ salt used here, so carry out this preparation well in advance of the rest of the experiment. Weigh out about 3.0 g of ferric ammonium sulfate (FeNH4 (SO4 ) 2 •12H2 O). Use a mortar and pestle to grind the salt to a fine powder. Accurately weigh 2.41 g of the powder into a 100 mL beaker and add 20 mL of concentrated sulfuric acid (see safety notes). Leave powder to soak in acid overnight. The next day, carefully pour the acid/powder slurry into a 500 mL volumetric flask, rinsing the beaker into the flask a few times with water, then make up to the mark with distilled water. Let this solution stand for several days until the ferric ammonium sulfate powder has fully dissolved. If possible, insert a magnetic stirrer bar and stir the solution to speed up this dissolving process.
Preparation of 1 mol L−1 ammonium thiocyanate solution
Weigh 38 g of solid ammonium thiocyanate into a 500mL volumetric flask and make up to the mark with distilled water.
Preparation of 0.15 mol L−1 potassium permanganate solution
Only required for analysis of iron tablet. Weigh 2.4 g of solid potassium permanganate into a 100 mL volumetric flask and make up to the mark with distilled water.
Preparation of iron tablet for analysis
Place iron tablet in a 100 mL beaker and use a measuring cylinder to add 20 mL of 1 mol L−1 sulfuric acid. Allow the tablet’s coating to break down and its contents to dissolve. You may help this process by using a stirring rod to carefully crush the tablet and stir the solution. (NB: iron tablets sometimes contain filler materials that may not fully dissolve in acid).
Once the iron tablet is dissolved, add 0.15 mol L−1 potassium permanganate solution dropwise, swirling the beaker after each addition. Iron tablets usually contain ferrous sulfate, with iron present as Fe2+ ions. Since Fe2+ does not form a coloured complex with thiocyanate, permanganate ions are added to oxidise all the Fe2+ to form Fe3+ ions. For the first few drops of permanganate, the purple colour will disappear immediately upon addition to the iron solution; however, as further drops are added, the colour will begin to linger for a little longer. Stop adding potassium permanganate drops when the purple colour persists for several seconds after addition − usually no more than about 2 mL of 0.15 mol L−1 permanganate solution will be required.
Transfer the iron solution to a 250 mL volumetric flask, rinsing the beaker with distilled water a few times and transferring the washings to the volumetric flask. Make up to the mark with distilled water, stopper the flask and mix well.
Use a pipette to transfer 5 mL of iron solution to a 100 mL volumetric flask and make up to the mark with distilled water. This diluted solution will be used for colorimetric analysis.
Calculations
Using only the absorbance results obtained for your Fe3+
- standard solutions (not your unknown iron sample), prepare a graph with [Fe3+] (in mol L−1) as the horizontal axis and absorbance (at 490 nm) as the vertical axis.
- Draw a line of best fit for your data points that goes through the origin (because absorbance must be zero when Fe3+ concentration is zero).
- Now identify the point on your line of best fit which corresponds to the absorbance measured for your unknown iron sample. By drawing a vertical line to the horizontal axis you will be able to determine the concentration of Fe3+ in your unknown solution.
- Use this concentration to calculate the mass of iron (in mg) in your original tablet or food sample (NB: the molecular weight of iron is 55.8 g mol−1). Remember to take into account any dilutions that you performed while preparing your sample solution.
- If the absorbance value you measured for your unknown iron sample is greater than the absorbance value for your highest concentration Fe3+ tandard you will need to modify the above procedure. In the case of an iron tablet, you should repeat the analysis with a more dilute solution of the dissolved iron tablet. In the case of a food sample, you should repeat the analysis using a smaller mass of your food.