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Just for fun!
Did you ever bite on something really hard while eating TOTAL cereal?
High iron content cereals, such as TOTAL, have been known to give you all the iron you need (100% RDA). What the cereal companies didn’t tell you is that you are actually eating pure iron fillings! Try mixing TOTAL cereal with water and putting a magnet in it. Pull out the magnet and notice all the little flakes of iron fillings sticking to your magnet. Is this a problem when we eat the cereal? Are we able to digest this iron? Check out the following website for the answer:
Useless Information: TOTAL Cereal

What is in my vitamin pill besides ‘vitamins?
How do the vitamins retain those cool little shapes? Is the entire pill filled with the minerals and vitamins listed on the label? Most pills have approximately 50% waste materials in them, such as fillers (to take up space in the pill), binders (glue to hold the pill together) and other odds and ends, such as, lubricants, colorants, flavors (in chewable tablets) and plasticizers.
KareMor Manufacturing International


laughing smileyJokes
A chemistry graduate student had the fortune to share their space with a cat they happened to name Ion. The student loved to introduce their feline friend as their cation!

A physicist, biologist and a chemist were going to the ocean for the first time.
The physicist saw the ocean and was fascinated by the waves. He said he wanted to do some research on the fluid dynamics of the waves and walked into the ocean. Obviously, he was drowned and never returned.
The biologist said he wanted to do research on the flora and fauna inside the ocean and walked inside the ocean. He too, never returned.
The chemist waited for a long time and afterwards, wrote the observation, "The physicist and the biologist are soluble in ocean water."


Smiley face image taken from Microsoft Office Clip Art

Analysis of Ferrous Iron in a Vitamin Pill

Goals

  • Use the quantitative technique of spectrophotometry to determine the amount of iron in a vitamin pill.
  • Become comfortable with Beer's Law.
  • Practice dilution calculations.
  • To learn the process of error propagation through a multi-step calculation using lab data.

Background

Iron and the Human Body

The average adult human body contains 4 to 6 grams of iron.  The majority of this iron is found in the red blood cell protein called hemoglobin.  The function of hemoglobin is to transport oxygen from the lungs to the various tissues in the body where it is used to produce energy.  Humans obtain iron from their diet in foods such as meats and leafy green vegetables.  Dietary supplements of iron can be taken to alleviate and/or prevent anemia, a condition caused by lack of iron, which results in low energy and pale skin tone.  Most vitamin tablets contain iron in the form of ferrous fumarate, [Fe2+(C4H2O4)2-].

Spectrophotometric Analysis

Spectrophotometry is one of the most common techniques used in quantitative analysis of samples for a specific chemical component.  In spectrophotometry, the amount of electromagnetic radiation that a sample absorbs is measured with an instrument called a spectrophotometer.  This absorbance is related to the concentration of the analyte, the chemical compound being analyzed, by a relationship called Beer's Law:

A = εbc (equation 1)

Where: 

A = absorbance
ε = molar extinction coefficient (a.k.a. molar absorptivity)
(ε is a wavelength-dependent constant that is unique for each chemical species)
c = concentration
b = path length, typically 1.0 cm

Beer's law shows a linear relationship between absorbance and concentration. The linear relationship holds true for absorbance values between 0.1 and 1.0.  Therefore, samples and standards must be prepared so that the concentration is within the linear range of the absorbance.

Standard Curve
Figure 1:  Beer's Law "standard curve"

In order to determine molar extinction coefficient, ε, you must make a calibration curve using a series of standard solutions (see Figure 1). A standard solution is a solution in which the concentration of the analyte is known.  The absorbances of several standard solutions are measured and absorbance vs. concentration is plotted.  The plot is called a "standard curve", and should be a straight line with y-intercept virtually zero, whose equation y = mx + b corresponds to the Beer's Law equation A = εbc. (Why should the y-intercept be zero?) In these equations, y corresponds to A, x corresponds to c, and the slope corresponds to εb. Once the calibration curve is complete, the absorbance of the unknown sample is measured, and the value of the extinction coefficient (ε) is determined in Beer's Law using the slope of the standard curve: εb. The Beer's Law formula can then be used to determine the concentration of the unknown.

Spectrophotometer

The spectrophotometer that we are using, the Ocean Optics Spectrophotometer, consists of a light source, a monochromator, a sample holder, and a detector.  The light source is a lamp that emits a broad range of wavelengths.  The monochromator separates the light into a narrow range of wavelengths that will reach the sample cell.  The detector measures the amount of light that gets through the sample, and compares it to the amount of light that passed through the blank sample. The result is given in units of %transmittance (%T) or absorbance (A), where A= -log (T).  We will use units of absorbance. Instructions for using the Ocean Optics Spectrophotometer are in the Appendix section of this Lab Manual.

When you use a spectrophotometer, you must first calibrate the instrument (also known as "blanking" or "zeroing"). To do this, you make a "blank" solution that is identical to the solution you will measure but contains none of the analyte that you wish to measure. You put the blank solution in a plastic cuvette in the spectrophotometer and calibrate the instrument to the blank, setting the absorbance to zero when the blank is present. Thus, when you put your solution containing analyte into the spectrophotometer, it will only measure absorbance due to the analyte.

Making solutions: Limiting reactants and reactants in excess

In this lab, you will be performing a number of reactions, including iron ions reacting with molecular phenanthroline ions to form an ionic iron-phenanthroline complex:

Fe2+ + 3 Phen→ [Fe(Phen)3]2+ 


Image from: http://water.me.vccs.edu/courses/env211/changes/ironreaction.gif

Check out the cool 3-D rendition of this huge complex!

This reaction equation indicates that you need three molecular PhenH+ ions to react with each Fe2+ ion to form one [Fe(Phen)3]2+ ion. Likewise, exactly 300 Phen molecules will react with exactly 100 Fe2+ ions to make exactly 100 [Fe(Phen)3]2+ ions. However, if exactly 310 Phen molecules are combined with exactly 100 Fe2+ ions, 100 [Fe(Phen)3]2+ ions will be formed and 10 Phen molecules will remain unreacted. We say that Phen molecules are "in excess" because there are excess unreacted Phen molecules in solution when the reaction reaches completion. Similarly, we say that Fe2+ ions are the "limiting reagent" because all of the Fe2+ ions are used up in the reaction and the amount of products formed are limited by the number of Fe2+ ions initially available.

Making solutions: Know your glassware

In order to make our standard curve, we need to know the concentration of our analyte (the iron-phenanthroline complex; see "Synopsis" below) precisely. This means that we need to know 1) the precise volume, and 2) the precise amount of analyte. We will be performing reactions to make our analyte. Since iron is our limiting reactant for all our solutions, it's quantity is the only one that needs to be measured precisely. All the other reactants are in excess, which means that there will be plenty of them available to react even if our amounts are inexact, so we do not need to measure the amounts precisely.

105 Glassware

Figure 3:  Glassware

If you need to know a volume precisely, you need to use volumetric glassware (see Figure 3). If you do not need to know a volume precisely, you can use a graduated cylinder (see Figure 3), which can save a lot of time.

Synopsis of the Experiment

In this experiment, you will determine the amount of iron in a vitamin pill, using spectrophotometry to measure absorbance, Beer's Law to calculate concentration, and stoichiometry to calculate milligrams.  The experiment will be carried out over the course of two labs.

Part 1. Making a Standard Curve

In the first part, you will make standard solutions of an iron-phenanthroline complex and measure their absorbance. You will start with a colorless solution of ferrous ammonium sulfate (Fe(NH4)2(SO4)2 • 6H2O), which dissociates in water to yield one Fe+2(aq) ion for each ferrous ammonium sulfate molecule. You will perform a reaction to turn colorless Fe+2(aq) into a colorful iron-phenanthroline complex which can be measured spectrophotometrically. The reaction is as follows:

First, you need to ensure that all of the iron is in the Fe+2(aq) form. Iron exists predominately in three forms: Fe(s), Fe2+ and Fe3+. Acid has already been added to the ferrous ammonium sulfate solution to ensure that any Fe(s) impurities are oxidized to Fe3+(aq). Your first step will be to add a reducing agent, hydroxylamine hydrochloride (HONH2•HCl), to ensure that any Fe3+(aq) impurities are converted to Fe2+(aq).

Next, you will add sodium acetate (NaOCOCH3), a buffering reagent, to the ferrous ammonium sulfate solution to make sure that the solution has a pH that is appropriate for the next step: formation of the spectrally active iron-phenanthroline complex. 

In the third step, you will react your Fe+2 solution with 1,10-phenanthroline (C12N2H8).  One Fe2+ ion forms a complex with three 1,10-phenanthroline molecules as seen in both reactions below:

Fe2+ + 3 Phen→ [Fe(Phen)3]2+ 

In solution, the iron-phenanthroline complex is a bright yellow-orange color, with an absorbance maximum at 508 nm (see Figure 4). You will measure the absorbance of your standard solutions. Plotting this data as A vs. c will result in a linear fit with slope εb, according to Beer’s Law.

absorbance spectrum for Fe(II) Phenanthroline Compex
Figure 4: Absorbance Spectrum for Fe(II) Phenanthroline Complex

Part 2. Making a Vitamin Solution and Measuring its Absorbance

In the second part of the experiment, you will crush a vitamin tablet and perform a similar series of reactions that will turn the iron from the pill into the colorful iron-phenanthroline complex whose concentration can be determined spectrophotometrically.

In order to convert the iron from the tablet into a spectrally active species (the iron-phenanthroline complex), we must start with Fe+2. To guarantee that all of your iron from the vitamin pill is in the Fe2+ state, you will add acid to oxidize any Fe(s) to Fe3+. Then you will add a reducing agent, hydroxylamine hydrochloride (HONH2•HCl) to reduce Fe3+ to Fe2+

Once again, sodium acetate (NaOCOCH3) is added to make sure the solution has a pH that is appropriate for the final step.  In this final step, you will react your Fe+2 solution with 1,10-phenanthroline (C12N2H8) to form the spectrally active iron-phenanthroline complex.  One Fe2+ ion forms a complex with three 1,10-phenanthroline molecules.

Fe2+ + 3 Phen→ [Fe(Phen)3]2+ 

Once you have created the iron-phenanthroline complex, you can determine the amount of iron in your vitamin pill by spectral measurements, using Beer’s Law and the extinction coefficient that you determined in the first lab of the experiment.

Preparation – Part 1

Reading Assignment:

Questions:

Fill out the prelab worksheet that can be found after the Experiment section for Part 1.

Experiment – Part 1

To print instructions, select the portion that you with to print, choose File/Print, and choose "selection" to prevent printing the entire document.

Safety

**Ferrous ammonium sulfate and sodium acetate solutions are mild irritants. Please rinse any affected area with water.

**0.002 M 1,10-Phenanthroline solution is toxic if ingested. It's harmful to the environment, so please dispose of this chemical in the appropriate waste container.

** 2% Hydroxylamine hydrochloride solution is harmful if inhaled, ingested, or absorbed through the skin. Possible mutagen. Please rinse any affected area with copious amounts of water.

**6M Hydrochloric acid (for Part 2) is a severe irritant to skin and eyes.
Please use gloves when working with this solution and change your gloves frequently. Please rinse any affected area immediately with copious amounts of water.
[ Please neutralize any hydrochloric acid spills with saturated sodium hydrogen carbonate solution.]

Do not pour solutions down the drain. Pour all chemical waste into the disposal containers provided.

Clean up after yourselves (You may be quizzed on this and will be expected to remember it for future labs):

  • Rinsed beakers and graduated cylinders go in the big orange bin by the door.
  • Volumetric flasks get rinsed and go back where they came from
  • Volumetric pipettes go tip-up in cylinder in front of room
  • Plastic cuvettes go in the trash
  • Test tubes and disposable pipettes go in the glass waste boxes
  • Disposable pipette bulbs are not disposable and go back where you got them.
  • Everything else gets rinsed and put back where you got it.

Materials and Equipment

  • 2% hydroxylamine hydrochloride, HONH2• HCl
  • 10% sodium acetate, NaOCOCH3
  • 5  x 10-4 M solution of ferrous ammonium sulfate in 1% sulfuric acid
  • 0.002 M 1,10-phenanthroline, C12N2H8
  • Ocean Optics spectrophotometer and laptop computer
  • 50.00 mL volumetric flasks
  • 1.00-8.00 mL volumetric pipets
  • 10, 25, and 100 mL graduated cylinders

Instructions

Students will work in groups of 2 unless otherwise specified by your instructor.

Along with other sets of lab partners, you will be making a set of 8 standard solutions (precisely known concentrations of iron-phenanthroline complex) plus a blank solution (no iron-phenanthroline complex ).  The standard iron-phenanthroline solutions will each contain a different amount of Fe2+ stock solution (measured using volumetric pipettes from 1.00 mL to 8.00 mL in volume). Follow your lab instructor’s directions for which samples to make:

  • Determine which standard solution(s) you and your lab partner will be making. Obtain one 50.00 mL volumetric flask for each solution you will be making.
  • You should never put a pipette into a reagent bottle. In order to avoid this, use a small beaker to obtain some ~5 x 10-4 M ferrous ammonium sulfate Fe2+ stock solution (10mL should be plenty). You can pipette from this beaker. Be sure to record the exact concentration of this solution, with the number of significant figures shown on its container.
  • For each of the standard Fe2+ solutions you will be making, pipette the appropriate aliquot of Fe2+ stock solution into a flask. If you are making a blank solution, don't add any Fe2+ stock solution, but keep an empty volumetric flask to which you will add the chemicals listed below.
  • To each of these flasks, add, in this order:
    • 5 mL of 2%(v/v) hydroxylamine hydrochloride solution and
    • 10 mL of 10%(w/v) sodium acetate solution
    • 25 mL of 1,10-phenanthroline solution
    • Fill to the mark with deionized water.  Mix well.

You and your lab partner will make your own calibration curve using all eight of the Fe2+ solutions prepared by everyone in the lab.  See the Appendix for directions on how to blank and appropriately use the spectrophotometer.  Calibrate the spectrophotometer with your blank solution..  Measure the absorbance at 508 nm for each of the iron(II) standards.

Make a plot of absorbance as a function of iron-phenanthroline complex concentration (M, not mL).  Use (0,0) as a data point (this is your blank), but do not choose "Set intercept =" in the Add Trendline dialog box (Options tab).  Have the Excel program print out the equation for the line (in the form of y = mx + b) and the regression coefficient (R2) on your plot. Before you leave, show your plot to your instructor and enter your extinction coefficient in the class spreadsheet. You will use the class average extinction coefficient for your calculations next lab.

Prelab & Results – Part 1

Click here for Prelab worksheet for Part 1: pdf or Word format

There is no Results worksheet for Part 1.

Preparation – Part 2

Reading Assignment:

  • Description of Experiment—see below
  • Read the paragraphs below on precision, accuracy, and uncertainty

Print:

Print and bring to lab Workshop 3 in pdf or Word format. Please print double-sided. If you think it would be beneficial to you, feel free to attempt the workshop before you come to lab.

I. Precision and Accuracy

There is uncertainty (a.k.a. "error") in any measurement that you make. Both precision and accuracy errors lead to uncertainty:

  1. Precision (reproducibility error): A measurement is precise if repeated measurements yield very similar values. Precision is often limited by equipment, e.g. the number of significant figures available for a given instrument or piece of glassware.
  2. Accuracy (systematic error): A measure of how close measurements are to the correct value. The average of several accurate measurement will yield the correct value.

Accuracy Precision

Example: Let's say you're measuring the diameter of a penny using the ruler shown below. The smallest increments on the ruler are millimeters. Precision: When you take a measurement, only your last digit should have uncertainty in it. So you should be able to record a diameter to the nearest tenth of a millimeter, with uncertainty in the last digit. Your precision error would be between ± 0.1mm and ± 0.4mm, based on how well you think you can estimate that last digit. Accuracy: Your accuracy error, may be larger than your precision error if you measure from the end of the ruler, because the markings on the ruler begin 1mm from the end of the ruler. That would give you a systematic (accuracy) error of 1mm.

ruler

II. Uncertainty

Uncertainty can be determined in two different ways:

  1. Statistical uncertainty: When several repeated measurements are made, the standard deviation of the measurements can be taken as the total uncertainty.
  2. Measured uncertainty: The experimenter can determine the uncertainty for each measurement taken. The total measured uncertainty is the sum of the precision and accuracy errors. If many measured values are combined to calculate a final value, the uncertainty of the final value can be calculated using propagation of error (see below).

Uncertainty should always be reported to only one significant figure.

There are two ways to represent uncertainty:

  1. Absolute uncertainty (AU) is the uncertainty in a measurement, given in the same units as the reported value. For example, the grape's width is 12.3 ± 0.2 mm, where 0.2 mm is the AU.
  2. Relative uncertainty (RU) represents AU as a fraction or percentage. RU = AU/|value|.
    For example, 0.2mm/12.3mm = 0.02 = 2%. The grape's width is 12.3 mm ± 2%, where 2% is the RU.

Propagation of error refers to the fact that if you perform calculations with values that each have an uncertainty associated with them, you can follow a simple set of rules to determine the uncertainty in your final calculated value.

There are a few rules that govern how uncertainty propagates throughout a calculations:

  • Addition and Subtraction: AU = sum of AU's

    When calculating uncertainty for the sum or difference of measured values, AU of the calculated value is the sum of the absolute uncertainties of the numbers being added and subtracted. 
  • Example:

    A = 79 ± 8 (AU = 8)

    B = 28 ± 4 (AU = 4)

    AU of (A + B) = 8 + 4 = 12 = 10

    A + B = 107 ± 12 = 110 ±10 (I have used underlines to indicate sig figs)

    Notes:

    • To determine the correct number of significant figures in your answer, calculate the absolute uncertainty, round it to one significant figure, and then round the calculated value to the same digit.
    • RU can be calculated using the equation RU = AU/|value|.
    • Even if you are subtracting measured values, be sure to add AUs.

  • Multiplication and Division: RU = sum of RU's

    When calculating uncertainty for the product or ratio of measured values, RU of the calculated value is the sum of the relative uncertainties of the individual terms. 

    Example:

    A = 79 ±8

    B = 28 ±4

    AB = 2212

    RU of AB = RUA + RUB = 8/79 + 4/28 = 0.2441 = 0.2

    RU = AU/|value|, so AU = RU*|value|

    AU of AB = RU*|AB| = 0.2441*|2212| = 540 = 500

    AB = 2212 ± 540 = 2200 ± 500

    Notes:

    • Uncertainties are always positive.  Be sure to calculate RU using absolute values.

  • Combination of addition/subtraction with multiplication/division
  • For a combination of these operations we follow the standard order of operations. 

    Example:

    A = 79 ±8

    B = 28 ±4

    AB + A = 2291

    RU of AB = RUA + RUB = 8/79 + 4/28 = 0.2441 = 0.2

    AU of AB = RU*|AB| = 0.2441*|2212| = 540 = 500

    AU of (AB + A) = AUAB + AUA = 540 + 8 = 548

    AB + A = 2291 ± 548 = 2300 ± 500

Refer to Notes on Uncertainty Analysis in the lab manual appendix for more examples and a more complete discussion of uncertainty propagation. Refer to Summary of Significant Figures and Uncertainty in the lab manual appendix for a summary of the rules of uncertainty propagation.

Questions:

Fill out the prelab worksheet that can be found after the Experiment section for Part 2.

Print:

Print and bring to lab (please print double-sided):

  • Workshop 3 in pdf or Word format
  • Uncertainty practice worksheet: pdf or Word format

Experiment – Part 2

To print instructions, select the portion that you with to print, choose File/Print, and choose "selection" to prevent printing the entire document.

Safety

Do not pour solutions down the drain. Pour all chemical waste into the disposal containers provided.

Materials and Equipment

  • 6 M HCl
  • 2% hydroxylamine hydrochloride, HONH2• HCl
  • 10% sodium acetate, NaOCOCH3
  • 0.002 M 1,10-phenanthroline, C12N2H8
  • vitamin tablets
  • Ocean Optics spectrophotometer and laptop computer
  • watch glass
  • fluted filter paper
  • 50.00 mL & 100.00 mL volumetric flasks
  • 1.00-8.00 mL volumetric pipets
  • 10, 25, and 100 mL graduated cylinders

Instructions

Each student will analyze her own vitamin. 

Preparation of the sample solution

Record the full brand name of the vitamin that you take and the mg of iron stated on the bottle. Take your assigned vitamin pill and grind it with a mortar and pestle. Transfer the powdered pill into a 250 mL beaker. Measure 25 mL of 6M HCl.  Using a disposable pipet, rinse the mortar with small volumes of your HCl solution; this will insure quantitative transfer of the powder. After you have rinsed the mortar, add the remainder of the 25 mL of HCl to your beaker. Cover the beaker with a watch glass.   Let it boil gently under the fume hood for 15 minutes.  Do not let the solution boil dry!  If the volume falls below about 15 mL, add deionized water.  Do not assume that an unlabeled container contains deionized water.  Let the solution cool to the touch in the hood.

Take your vitamin solution and gravity filter out any solids, using a a 100.00 mL flask, a long-stem funnel and fluted filter paper.  Filter into the 100.00 mL volumetric flask and, when it's done filtering, fill to the mark with deionized water. Mix well. Pipet 1.00 mL of this solution into a 50.00 mL volumetric flask and treat it as the standards were treated in Part 1 (add hydoxylamine hydrochloride, sodium acetate and 1,10-phenanthroline (in that order) and fill to the line with deionized water).  Before you measure the absorbance of this solution, make a blank according to the instructions below.

Some of the vitamins may be colored with dyes that absorb at 508 nm.  We can correct for this by making a blank that includes vitamin solution but is not reacted to form the iron-phenanthroline complex.  To make this blank solution, take a 50.00 mL volumetric flask and add 1.00 mL of your vitamin pill stock solution (from the 100.00 mL flask). Add the usual amounts of hydoxylamine hydrochloride and sodium acetate (in that order) and fill to the line with deionized water. Do not add 1,10-phenanthroline, because we do not want the iron-phenanthroline complex to form! Use this blank solution to zero your spectrophotometer. Then measure the absorbance of your vitamin pill solution with the 1,10-phenanthroline to get the amount of absorbance that is due to only the phenanthroline complex. 

For your results, you will be using your class's average extinction coefficient for the iron-phenanthroline complex. This value, along with the absorbance of the phenanthroline complex from your vitamin, will be used to calculate the concentration of your vitamin pill solution. Finally, you will use the concentration to calculate the amount of iron in your vitamin pill. Don't forget that you diluted your original 100.00 mL of vitamin solution before you measured its absorbance!

Workshop 3: Uncertainty

Click here to download Workshop 3: Uncertainty in pdf or Word format

After completing the workshop, do the Uncertainty Analysis Practice worksheet that you printed before coming to lab (pdf or Word format). Check your answers here: pdf or Word format

Prelab & Results – Part 2

Click here for Prelab worksheet for Part 2: pdf or Word format

Click here for Results worksheet for Part 2: pdf or Word format

You should be prepared to do uncertainty propagation calculations for your next quiz.

Next Lab: Lab Practical

References

Laura Muller, “Spectrophotometric Determination of Iron in a Vitamin Tablet,” General Chemistry Lab, Wheaton College, 2000.

For more information:
MEDLINEplus Medical Encyclopedia: Serum iron
http://www.nlm.nih.gov/medlineplus/ency/article/003488.htm

Lab Tests Online: Iron Tests:
http://www.labtestsonline.org/understanding/analytes/iron/overview.html
http://www.labtestsonline.org/understanding/analytes/serum_iron/test.html

The method used in clinical chemistry labs for determination of serum iron:
Iron Panel of the International Committee for Standardization in Hematology.  Revised recommendations for the measurements of the serum iron in human blood. British Journal of Hematology 75:615-616, 1990.
Iron Panel of the International Committee for Standardization in Hematology.  Recommendations for the measurements of the serum iron in human blood. British Journal of Hematology 38:291-294, 1978.

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