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2617 Quantitative Analysis of Iron

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    1. Objective

    After completing this experiment, the student will be able to analyze the iron content in a sample using spectrophotometry.

    1.2 Background

    Common iron supplement tablets contain various amounts of Fe as ferrous sulfate. One example is 65 mg Fe in 325 mg ferrous sulfate. How reliable is this information? This experiment will enable you to check up on the manufacturer.

    In this experiment, the presence of iron(II) ion will be detected by reacting the sample with 1,10-phenanthroline (Equation 1).


    (Equation 1)

    Three molecules of 1,10-phenanthroline will chelate with iron(II) to form a tris-1,10-phenanthrolineiron(II) orange-red complex. To ensure fast and optimal color development, this reaction will be maintained between a pH of 2.9 and 3.3 in the presence of excess phenanthroline. It is important to note that interferences to this reaction can come from strong oxidizing agents, phosphates, and other metal ions such as chromium, zinc, cobalt, copper, and nickel. The errors caused by these interferences can be minimized by using a larger excess of phenanthroline, initial boiling with acid, and adding excess hydroxylamine.

    Spectroscopy is one of the most powerful analytical techniques in modern science. Before the advent of spectrophotometric techniques, a chemist interested in determining the amount of a particular substance present in a sample had to analyze the sample via a series of chemical reactions specific to that species and then carefully weigh the products (entire tomes exist detailing such analytical reactions). This process was extremely time consuming, prone to error, and generally impractical for measuring trace amounts. Today, most routine assaying is done quickly and efficiently by means of spectroscopy.

    Spectroscopy works by correlating the concentration of a species in solution to the amount of light it absorbs. In this experiment, we will determine the quantity of iron (as Fe-phen) using visible absorption spectroscopy. Because the wavelengths of light we will use are in the visible portion of the electromagnetic spectrum, our solutions will all be colored. However, this technique can also be used in other regions of the spectrum where the wavelength is not visible to our eyes but can be measured using a photodetector.

    1.3 Experimental Considerations

    1.3.1 Choosing a Wavelength

    It is important to choose a wavelength where the solution strongly absorbs light.

    The stronger the absorption at a particular wavelength the more sensitive the instrument will be at that wavelength and the more accurate your results.

    In practice, one first determines the best range of wavelengths at which to test a solution and then selects one "best" wavelength for the experiment by measuring the absorption of the solution in question across this range and choosing the wavelength that gives the greatest absorption. For example, for the orange-red complex described above, the experimenter might choose to measure the absorption of the sample at a range of wavelengths between 480 nm and 560 nm, and then narrow-in on the wavelength that gives the greatest numeric value of the absorbance.

    Please note that while some spectrometers (like Spec 200 that will be used in this experiment) can do this automatically, others require changing the wavelength manually and each time the wavelength is adjusted these instruments need to be re-zeroed or they will not work properly.

    1.3.2 Calibrating the Instrument

    Once the specific wavelength that will be used for the experiment is chosen the experimenter needs to calibrate his or her spectrometer so that the absorbance readings can be converted into useful data such as molarity (M) of the sample.

    The absorbance of light (A) is directly proportional to the concentration of the species in solution (C). This relationship is known as "Beer's law" and may be expressed as:

    A = εbC (Equation 2)

    where ε is the molar absorptivity (or the molar extinction coefficient), b is the path length or distance the light travels through the sample, and C is the concentration of the solution in terms of molarity. In most instruments b is a constant (1.0 cm in our experiment) and can therefore be factored into ε.

    Thus, Equation 2 can be simplified as:

    A = εC y=mx (Equation 3)

    which is the equation of a line where the intercept is through the origin (0,0).

    The molar absorptivity, ε, can therefore be determined by finding the slope of a plot of the Absorbance (A) as a function of concentration (C) for a series of standard solutions of known concentrations. This is known as a calibration curve.

    Once ε has been determined from the calibration curve, Equation 3 can be used to determine the concentration of an unknown solution by measuring its absorbance under

    the same conditions.

    N.B. A new calibration curve is required if a different instrument, wavelength, type of solution, or procedure is used. It is also a good idea to check an instrument's calibration on a periodic basis.

    Overall, this experimental project consists of:

    • Carefully preparing five (5) iron(II) solutions of known concentrations
    • Determining the wavelength at which the solution is absorbing the strongest, λmax
    • Measuring the absorbance of each known solution to prepare a standard calibration curve
    • Carefully preparing an iron(II) solution from a given dietary supplement
    • Calculating the amount of iron in one tablet (from the standard calibration curve) and comparing it with the manufacturer’s label

    Further Reading

    Technique I: Use of Spectrophotometer (from our laboratory manual)

    Technique E: Volumetric Transfer of Reagent Using a Pipet

    Expt 2501 Using Excel for Graphical Analysis of Data


    3.0 CHEMICALS AND SolutionS



    Approximate Amount


    Laboratory water


    500 mL


    Hydroxylamine solution

    2% w/v

    25 mL




    0.2 g

    Light green powder



    50 mL


    Sodium acetate

    10% w/v

    100 mL


    Hydrochloric acid

    6 M

    25 mL


    Sulfuric acid

    6 M

    6 mL


    Iron supplement tablet




    250 mL volumetric flask

    100 mL volumetric flasks (7 needed)

    500 mL volumetric flask

    250 mL beaker (2 needed)

    Volumetric pipets

    Graduated cylinder

    600 mL beaker (waste container)

    Filter paper


    Watch glass

    Spec 20/200



    1. Using an Iron Supplement Tablet to Prepare a Sample Solution

    NOTE: If unfamiliar, ask your instructor to describe and demonstrate the use of volumetric pipets and volumetric flasks.

    1. Read the label of the bottle of iron supplement to find the amount of Fe in each tablet. Record these values in section 6.4 on your data collection sheet.
    2. Transfer the tablet into a 250-mL beaker. Add 25 mL of 6 M HCl into the beaker. Check the pH of the solution using a pH paper. This pH should be around 2.
    3. Boil the solution gently for about 15 mins on a hot plate to COMPLETELY dissolve the solids. Do not let the volume of the mixture to fall below 15 mL, in which case, add some laboratory water. Turn off heat and let it cool.
    4. Gravity filtration: Filter the solution into a clean volumetric flask. (Note: Use 100-mL capacity if the iron supplement tablet contains 28 mg or less; 250-mL capacity if the iron supplement tablet contains 65 mg.) Make sure no particles of dust, paper, or other impurities get into the flask. Pour the iron-containing liquid from the beaker onto the filter paper in the funnel. The clear solution that comes through the filter contains your sample. The solid that is left on the filter paper will eventually be discarded. Rinse the beaker three times with laboratory water (about 30 mL total) and pour these rinsings into the filter funnel also.
    5. When the sample has been filtered, remove the funnel, and dispose in the trash the filter paper and undissolved residue from the iron supplement tablet.
    6. Carefully add enough laboratory water to bring the solution level exactly to the mark on the flask. Mix the sample thoroughly by carefully inverting the stoppered flask repeatedly. (Hold the glass or rubber stopper tightly in place with your thumb. Make sure no liquid leaks out before the solution is completely mixed.)
    7. The solution at this point may be too concentrated to be used. Prepare a dilution of it by using a volumetric pipet to transfer exactly 1.00 mL of it into a clean 100.0 mL volumetric flask. (The flask may be wet as long as it is clean.) Carefully add 10.0 mL of sodium acetate, 2.0 mL hydroxylamine, and 4.0 mL phenanthroline in this order. Add laboratory water to bring the solution level up to the 100.0-mL mark on the flask. Mix thoroughly. This will be your sample solution. You may label this as ‘Sample #1’. Set aside on your lab bench.
    8. Repeat this preparation for another iron supplement tablet as an additional trial. Depending on time constraints, students may perform a total of three trials, or they may borrow data from another lab group for statistical comparisons. Make sure the other lab group is using the same tablet brand as your group!

    5.2 Preparation of Iron(II) Stock Solution

    1. Into a 250 mL beaker containing around 10 mL of laboratory water, slowly add 6 mL of 6 M sulfuric acid.
    2. Weigh approximately 0.140 g of Fe(NH4)2(SO4)2.6H2O (record its exact mass) and transfer into the beaker containing the acid. Mix thoroughly to dissolve the solid completely.
    3. Transfer the solution into a 500 mL volumetric flask. Rinse the beaker three times with laboratory water (about 30 mL total) and pour these rinses into the volumetric flask. Add laboratory water exactly to the mark. Mix the sample thoroughly. Label this as ‘Iron(II) stock solution’.
    4. Calculate the exact concentration of this iron(II) stock solution and record on your data recording sheet.

    5.3 Preparation of Solutions of Known Concentration

    1. Obtain the following solutions and transfer into small beakers: 75 mL iron(II) stock solution, 100 mL sodium acetate, 25 mL hydroxylamine, and 50 mL phenanthroline.
    2. Label five (S1, S2, S3, S4, S5) clean 100 mL volumetric flasks for the preparation of your standard solutions.
    3. Using a volumetric pipet, carefully measure 1.00 mL of the stock solution you made in section 5.2 and place it in a volumetric flask (S1). Add mL 10.0 mL sodium acetate, 2.0 mL hydroxylamine, and 4.0 mL phenanthroline (in this order). Carefully fill the flask with laboratory water exactly to the 100 mL mark. Mix the sample thoroughly. Allow at least 5 minutes for maximum color development. This is your standard solution S1.
    4. Repeat step 3 to prepare the other standard solutions - using 2.00, 3.00, 5.00, and 10.00 mL of the stock solution from section 5.4 into each volumetric flask. Calculate the concentrations of these standard solutions and record on your data recording sheet.
    5. Prepare a blank solution (labeled S0) by combining 10.0 mL sodium acetate, 2.0 mL hydroxylamine, and 4.0 mL phenanthroline in a clean 100.0 mL volumetric flask and diluting with laboratory water to the mark. Mix thoroughly.

    5.4 Absorbance Measurements

    1. Gather all the prepared solutions (standards and sample) for spectroscopic analysis. The proper use and operation of the spectrophotometer can be found in Appendix: Technique I Use of Spec 20/Spec 200.
    2. Obtain solution from S5 to obtain the λmax that will be used for the determinations. Use the ‘Scan’ mode of the spectrophotometer.
    3. Once λmax is determined, adjust the spectrophotometer to measure the absorbance of the solutions at this wavelength.
    4. Use the ‘blank’ solution (S0) as the reference liquid - that is, set the instrument for zero absorbance when the light path is passing through this ‘blank’ solution.
    5. Measure the absorbance for each of your solutions of known concentrations (S1- S5) and for your sample solutions (the 100 mL volumetric flasks labeled "Sample#").
    6. Rinse the cuvette (sample holder) thoroughly with laboratory water, then with each solution in between measurements.


    Last Name

    First Name


    Partner Name(s)



    6.1 Preparation of the Iron(II) Stock Solution (procedure section 5.2)

    Mass of Fe(NH4)2(SO4)2.6H2O (green powder)


    Concentration of Fe2+ stock solution


    Show your calculations below for the exact concentration of your iron(II) stock solution.

    6.2 Absorbance Measurements (procedure section 5.4)

    λmax = ______________________________

    6.3 Absorbances for Standard Solutions














    Show your calculations below for one of the standard solutions.

    6.4 Absorbances for Tablets


    Mass of Fe

    (According to Manufacturer)


    Tablet #1


    Tablet #2


    Tablet #3

    (or class data)



    1. Using Excel, construct a graph of absorbance (A) vs. concentration (C) for your standard solutions. Provide a comprehensive title for the graph and label your axes completely. Using the trendline function, add a best-fit line to your plotted data and display the equation of this line (in the form: y = mx + b or A = εC + b) and its R2 value on your graph. Submit this graph with your report.

    Equation of the line_____________________________________________

    Use this graph to determine the value of the molar absorptivity, ε. (Be certain to include the proper units!)

    Molar absorptivity, ε: ______________________________

    2 Use your graph to determine the concentration and mass of Fe2+ in your tablets.


    Tablet #1

    Tablet #2

    Tablet #3 or class data



    Concentration of Fe2+in the diluted solution (from calibration curve)


    Concentration of Fe2+(aq) in the original undiluted sample


    Mass of Fe in the tablet*


    Average mass of Fe in the tablet


    *use the undiluted original concentration

    Show your calculations for one of the trials.

    1. Discuss your experimental results. Do your results agree with the iron supplement manufacturer's claim regarding the milligrams of Fe in each tablet? Discuss any of your experimental errors that may affect your conclusions.

    1. How does the structure of the Fe-phen complex become more useful in the assay of iron using spectrophotometry. Consider the color and the stability of the complex.

    1. Using the analytical method we used in this lab, what other types of samples could you assay the Fe content of? Discuss any limitations or interferences with the assay process that you may encounter. How would you plan to address them?

    2617 Quantitative Analysis of Iron is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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