Chemistry Lab 1 - Spectrophotometric Determination of Iron in Dietary Tablets*
January 9, 2001

* Adapted from "Exploring Chemical Analysis", by D. C. Harris, Freeman, 1997.
 

Introduction. In this lab you will perform a colorimetric analysis to determine the amount of iron in a vitamin pill. As the name implies, colorimetric analyses rely on the color, specifically the intensity of the color, of a solution to determine the concentration of a particular component. For example, common sense and experience tell you that the depth of color of solutions of food coloring vary depending on how much dye is present. As it turns out, the relationship between color and concentration is given by Beer's Law, as explained below.

When visible light with an initial intensity P_o passes through a colored medium, some of the light is absorbed and the power is diminished to P. The transmittance, T, of light is the ratio of the power of the light that emerges from the sample to that of the incident beam.
 
 

Another way of expressing how much light is absorbed by a sample is its absorbance, A. Absorbance is related to transmittance in the following way:

Since absorbance is a logarithmic quantity, a change of one absorbance unit represents an order of magnitude change in transmitted light. For example, if a sample transmits 10% of incident light (T = 0.1), its absorbance is 1.0. An absorbance of 2.0 represents a transmittance of 1%, an absorbance of 3 represents a transmittance of 0.1%, and so on.

The advantage in using absorbance is a measurement of light absorption is that it is directly proportional to concentration. This relationship is known as Beer's Law.

Here, b is the pathlength (in cm) of the sample, c is the concentration of the absorbing species (in mol/L) and e is the constant of proportionality (with units of L mol^-1 cm^-1) known as the molar absorptivity, or the extinction coefficient. The units of the molar absorptivity are such that absorbance is a dimensionless quantity, as its original definition dictates.

For any substance e is a function of wavelength, thus absorbance varies across the visible spectrum. Species that absorb blue light strongly appear yellow since they transmit mainly in the green, yellow and red portions of the spectrum. Similarly materials that absorb red light strongly will appear blue-green.

In this experiment, you will use Beer's Law to determine the iron content of a vitamin pill (Flintstones!). The compound that you will form is [Fe(phen)₃]²⁺, which is a deep red-orange color (l_max = 510 nm). Typically this type of colorimetric analysis entails the generation of a calibration curve, which is a graph of the absorbance of a series of solutions (standard solutions) against their concentrations. The slope of the resulting graph is equal to eb. Since the pathlength is known, the slope can be used to find the value of e and, thus, provides the needed information to use Beer's Law in the analysis of an unknown solution.
 

Procedure. Working in teams of two, follow the procedure as outlined below. Use an Ocean Optics fiber optic spectrometer to measure the absorbances of your iron solutions. Report the iron content of your tablet (in mg) to the rest of the class for statistical analysis.

1. Place one tablet in a 100 mL beaker and boil gently (in a fume hood) with 25 mL of HCl for 15 minutes. Filter the solution into a 100 mL volumetric flask. Wash the beaker and filter several times with distilled water to complete a quantitative transfer. Allow the solution to cool, dilute to the mark and mix well. Label this "Solution A".

2. Pipet 5.00 mL of Solution A to a fresh 100 mL volumetric flask and dilute to the mark with deionized water. Label this "Solution B". Set this aside until step 4.

3. Prepare 5 standard iron solutions by adding the following solutions to individual 100 mL volumetric flasks. Use volumetric pipets for all solutions except the sodium citrate, for which a measuring pipet can be used. Mix well and dilute to the mark with deionized water.
 
 
 

	Standard Number:
	1	2	3	4	5
Iron Standard (7.16 ´ 10-4 M)	0 mL	1.00 mL	2.00 mL	5.00 mL	10.00 mL
Phenanthroline	3.00 mL	3.00 mL	3.00 mL	3.00 mL	3.00 mL
Hydroquinone	2.00 mL	2.00 mL	2.00 mL	2.00 mL	2.00 mL
Sodium Citrate	0 mL	0.4 mL	0.8 mL	2 mL	4 mL

4. Pipet 10.00 mL of "Solution B" into a fresh volumetric pipet. Add 4 mL of sodium citrate, 2.00 mL of hyrdoquinone and 3.00 mL of phenanthroline solution. Dilute to the mark with deionized water and mix well. Label this "Solution C". Do not analyze solutions A or B.

5. After allowing the color of the solution to develop for 15 minutes, measure the absorbance of each of the standards and your unknown (solution C) at 510 nm. Use deionized water as a background in your measurements.

6. For the standard solutions, plot the measured absorbance as a function of iron concentration. From the slope and the known pathlength, calculate the molar absorptivity of the [Fe(phen)₃]²⁺. Use this information to calculate the concentration of iron in the analyzed solution. Finally, calculate the mass of iron (in mg) that was in the tablet you started with. Report this value to the rest of the class.

Report. A preliminary graph and your final calculations will be due on 1/16. Analyze the class data and present your results in a formal lab report due on 1/23. State what your result was and provide a statistical analysis of the class data (include mean, standard deviation and 95% confidence interval). Compare to the value that is stated on the label. Comment the accuracy of the result and suggest sources of any descrepancies. Include in your report a copy of your calibration curve.