Written by Arturo
Objectives: To determine the amount of iron in a vitamin tablet using visible spectrometry.
Many compounds form colored solutions. Potassium permanganate solutions are purple, many copper(II) solutions are blue, and most nickel(II) solutions are green. If a solution is colored, a species within it is absorbing part of light in the visible region. A graph of absorbance, the amount of light absorbed by a species at different wavelengths is called an absorption spectrum. A similar graph may be obtained showing the amount of light that is transmitted at various wavelengths, called a transmittance spectrum. Both absorption and transmittance spectra can be used to determine the wavelength at which an analyte absorbs most. Also, this wavelength is usually chosen in order to achieve the highest sensitivity in the spectral measurement.
Instruments used in measuring the transmittance or absorbance of a sample at a specified wavelength are called spectrometers. The detector of a spectrometer is used to measure the energy (or intensity) of light at a specified wavelength. Transmittance is defined as the intensity of light transmitted through a sample solution divided by that transmitted through a blank solution. Since the correlation between the concentration and transmittance of an analyte is not linear, transmittance is not commonly used in quantitative analysis. To simplify the measurement, the absorbance is mathematically defined as: Absorbance = – log(Transmittance). The linear relationship of the concentration and absorbance of an analyte , which is called Beer’s Law: A=εbc, is commonly used in a spectrometric determination. In Beer’s law, A is the absorbance, ε is the absorptivity, b is the light path length in the sample, and c is the concentration.
Metal ions are essential to health. For examples, iron is necessary in oxygen transport and for energy storage, copper and zinc are essential in many enzymes. Enzymes are catalysts that accelerate the rates of reactions in human bodies. Many people take vitamin pills with mineral additives to ensure sufficient amounts of these necessary metals. Vitamin tablets with iron added are quite common. The iron(II) present in vitamins is usually in the form of ferrous fumarate, FeC2H2O4. It is interesting to note that iron(III) is generally useless for living systems. Many areas have soils with high iron(III) content but little or no iron(II), and thus are nonsupportive of plant life, despite their high iron content.
In this experiment, the amount of iron in a vitamin tablet will be determined. Since iron(II) is only weakly colored (pale blue-green when uncomplexed in aqueous solution), the iron (II) will be reacted with 1,10 phenanthroline (C12H8N2) prior to analysis. The iron(II) phenanthroline complexes , [Fe(phen)3]2+ is highly colored (orange) and easy to detect. Sodium acetate (NaC2H3O2) is added to control the acidity of the solutions. Iron(II) is easily oxidized to iron(III), which does not form a colored complex with phenanthroline. Hydroxylamine hydrochloride, (NH2OH)HCl, a reducing agent, is added to ensure that all of the dissolved iron remains in the form of iron(II) as specified in the reaction below:
2NH2OH∙HCl + 4Fe3+ + 5H2O = N2O + 4Fe2+ + 6H3O+ + 2Cl- Experimental procedure
Weigh an iron-enriched vitamin tablet (such as One-A-Day Plus Iron) to the nearest mg. Place the tablet in a mortar and pestle, and crush the tablet to a fine powder. Weigh precisely about 40 mg of the powder, transfer it to a 10 mL Erlenmeyer flask, add 5 mL of 0.01 M HCl, and place a magnetic stirring bar in the flask.
Place the Erlenmeyer flask in a 50 mL beaker, which is used as a water bath and contains about 20 mL water and a boiling chip. Place the beaker which contains the sample flask on a magnet-stirring hot plate. Boil the water bath and the sample solution with stirring for 20 minutes. Allow the solution to cool to room temperature. Depending on the type of vitamin used, much of the solid may not dissolve in the acid— this is normal. The remaining solid consists of binder and other inert ingredients. The iron will have all dissolved.
Filter the sample solution through a filtering disk in a Hirsch funnel using suction filtration, and transfer the liquid filtrate to a 10 mL volumetric flask (using a Pasteur pipet). Rinse the Erlenmeyer with 1 mL of deionized water. Filter the 1 mL rinse through the filtering disk. Save the rinse. Repeat with a second 1 mL rinse.
Transfer the combined rinses to the same 10 mL volumetric flask. Dilute to the mark with deionized water. This solution will be used as the “vitamin solution” in this experiment.
Make up the following four standards, one blank, and two unknown trials (1.00 and 2.00 mL of the vitamin solution) in seven separate 10 mL volumetric flasks.
Standard 1 Standard 2 Standard 3 Standard 4 Blank
Vol. of soln. A, mL 0.50
1.00 1.50 2.00 0.00 0.00 0.00
Vol. of soln. B, mL Vol. of soln. C, mL 1.00 1.00
1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00
Vol. of soln. D, mL 1.00
1.00 1.00 1.00 1.00 1.00 1.00
Trial 1 (1.00 mL vit. Soln.) Trial 2 (2.00mL vit. Soln.)
Fill each volumetric flask to the
Use the blank solution to adjust a spectrometer to “zero absorbance” at
Then, measure and record the absorbance readings for the four standard solutions and the two sample solutions, separately.
mark with deionized water. Stopper and
mix each solution thoroughly. the wavelength of 508 nm.
Calculate the exact concentration of Fe(II) in each of the four standards using dilution calculation.
Construct a Beer’s law plot using the absorbance readings of the four standards versus their exact concentrations. Because the spectrometer was zeroed with the blank solution, it is appropriate to force the Beer’s law plot through the origin.
From the Beer’s law plot or from the Beer’s law equation for the regression line, determine the concentration of iron(II) in the sample solutions.
Since the volume and concentration of the sample solutions are now known, the mass of iron in the sample solutions can be readily calculated. The two sample solutions respectively contained 1.00 mL and 2.00 mL of the original 10.00 mL vitamin solution prepared from vitamin powder. The mass of Fe(II) in the 10.00 mL original vitamin solution can also be calculated. The average of the two trials should be used.
The vitamin solution was prepared from a weighed subsample of the whole tablet. The total mass of Fe(II) in the whole tablet can be determined.