Experiment 2: Spectrophotometric Determination of Iron in Vitamin Tablets (Adapted from Daniel C. Harris’ Quantitative Chemical Analysis and R. C. Atkins, Journal of Chemical Education 1975, 52, 550.)
Experimental work to be done on February 24 + one hour scheduled on your own Notebook due on March 4 (by 4:00 pm ⇒ 20% late penalty each 24 hour period thereafter)
In this experiment, you will dissolve the iron in a vitamin supplement tablet, digesting the cellulosic matrix in the process, and then reduce the iron to Fe2+ with hydroquinone:
While freshly-dissolved aqueous Fe2+ is nearly colorless, we can impart an intense red color by a stoichiometric reaction of Fe2+ with three molecules of the ligand 1,10-phenanthroline (phen):
The absorption spectrum of the complex, often written as Fe(phen)32+, has a maximum at about 510 nm. This complex is stable indefinitely at pH values of 3 or higher. Measuring the analyte solution's absorbance at λmax is a sensitive method for determining iron concentrations. You will prepare a series of standard solutions containing known concentrations of Fe(phen)32+, as well as a solution with Fe from a vitamin tablet, and measure their absorbances on the Chemistry Department’s Beckman DU7400 spectrophotometer. Construction of a calibration curve using your standard solutions will allow you to determine both the molar absorptivity of the Fe(phen)32+ complex and, with a pair of measurements of the iron tablet solution's absorbance under the same conditions, the mass of iron that was present in your vitamin tablet.
All solutions from this experiment can go down the drain.
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• The hot HCl you are using in this experiment will release corrosive fumes, and will rapidly eat into any organic material with which it comes in contact. Wear gloves while working around the hot acid, and heat it only in a fume hood. Of course, wear safety glasses! • The iron stock solution contains dilute sulfuric acid (H2SO4). Unlike HCl, which will evaporate as a gas, sulfuric acid is not volatile. Even a small drop of dilute H2SO4, left out in dry air, will turn into a small speck of much more concentrated sulfuric acid. So clean up any spill with a paper towel; the paper towels used can be safely thrown in the trash. (The H2SO4 reacts with the paper, slowly burning it, and neutralizing the acid in the process.)
A. Basics 1. The stock solutions required for this experiment will be provided for you, having been lovingly prepared by your lab instructor and the lab preceptors. They will be as follows: ▪ hydroquinone: an aqueous solution containing 2 g/L, stored in amber bottles ▪ sodium citrate: 50 g/L in water ▪ 1,10-phenanthroline: 1.0 g/L in water, stored in amber bottles ▪ stock Fe2+ (nominally 0.04 mg Fe/mL): Prepared by dissolving 0.5602 g of reagent-grade Fe(NH4)2(SO4)2·6H2O in a 2 L volumetric flask containing 2 mL of 98%mass H2SO4 and diluted to the mark with DI water. 2. You must not contaminate or dilute the stock solutions everyone in the lab is relying on! When pipetting, pour out small portions of the reagent you need from the stock bottles into a beaker. Never insert a pipet or other glassware into a stock bottle, and never pour unused reagents back into stock bottles. 3. Regarding pipetting, ▪ Practice and verify your pipetting technique, using one of your 10 mL Class A volumetric transfer pipets and your green (10 ml) pipet pump. Confirm that you can reliably transfer 9.98 ± 0.03 g of DI water into a tared beaker on a top loading balance. Do not use an analytical balance for this. ▪ Use the accurate volumetric transfer (not graduated) pipets for all iron-containing aliquots. The volumes of the other solutions need not be measured very accurately. ▪ If you are having trouble with pipetting, please note the following tips:...
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