In the Laboratory
Analysis of Zinc Tablets: An Extension to a Stoichiometry Experiment
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Steven Murov* and Brian Stedjee Science, Mathematics and Engineering Division, Modesto Junior College, Modesto, CA 95350; *
[email protected] The addition of an analysis of zinc ion tablets to a previously published stoichiometry determination (1, 2) adds an interesting motivational aspect to this discovery-oriented experiment. It is possible to write at least four chemical equations for the titration of ferrocyanide solutions with zinc ion. K4Fe(CN)6(aq) + ZnSO4(aq) = K2ZnFe(CN)6(s) + K2SO4(aq)
(1)
K4Fe(CN)6(aq) + 2ZnSO4(aq) = Zn2Fe(CN)6(s) + 2K2SO4(aq)
(2)
2K4Fe(CN)6(aq) + ZnSO4(aq) = K6Zn[Fe(CN)6]2(s) + K2SO4(aq) (3) 2K4Fe(CN)6(aq) + 3ZnSO4(aq) = K2Zn3[Fe(CN)6]2(s) + 3K2SO4(aq)(4)
Before performing the experiment, students usually expect the reaction to be the standard double replacement (reaction 2). However, determination of the mole ratio of zinc ion to ferrocyanide by titration reveals that reaction 4 is consistent with the results. Once the stoichiometry is determined, the technique can be used to determine the concentration of zinc ion or ferrocyanide in solutions of unknown concentration or the amount of zinc ion in commercial zinc tablets. Experimental Overview, Results, and Discussion Students should prepare 250 mL of 0.025 M potassium ferrocyanide from K4Fe(CN)6 ⭈ 3H2O, and 250 mL of 0.050 M zinc sulfate from ZnSO4 ⭈ 7H2O. Twenty-five milliliters of the K4Fe(CN)6 solution is pipetted into a 250-mL Erlenmeyer flask and 5 mL of 3 M sulfuric acid, 4 drops of diphenylamine indicator, and 3 drops of freshly prepared 1% potassium ferricyanide (which oxidizes the diphenylamine to its colored form in the absence of ferrocyanide) are added. The solution is titrated with the zinc solution from a milky white appearance to a blue endpoint. After students have obtained three consistent titrations and determined the stoichiometry (within experimental error, the titration gives a zinc-to-ferrocyanide mole ratio of 1.50, in strong support of reaction 4), zinc ion or ferrocyanide ion solutions of unknown concentrations or zinc tablets can be analyzed. For the zinc tablet analysis, the students should grind a commercial 50-mg (according to the label) zinc tablet (we have successfully used the Vitasmart brand from K-Mart and Spring Valley from Wal-Mart) and add the solid to the 25 mL of potassium ferrocyanide solution before beginning the titration. To avoid over-titration, students should be alerted that this titration will take only a few milliliters, as the zinc tablet uses up most of the ferrocyanide and the titration
requires only about 3 mL. The original ferrocyanide solution should be at least 0.025 M or the zinc could consume all of it. Our results for the “50-mg” zinc tablets gave a satisfying value of 51 mg of zinc per tablet. Tablets containing 100 mg of zinc are also readily available, but either these have to be broken into halves or 50 mL of ferrocyanide should be used. Tablets containing other ions, multiple-vitamin pills, and zinc lozenges did not give satisfactory results in our studies. Interestingly, the amount of zinc in the tablet can be calculated without use of the mole ratio. Two hundred milliliters of the diphenylamine indicator is prepared by dissolving 2.0 g of diphenylamine in 100 mL of concentrated sulfuric acid. The diphenylamine in sulfuric acid is added to 100 g of ice in a 400-mL beaker with stirring. Because preparation of this solution is potentially hazardous, the indicator solution should be provided by the instructor. Hazards Diphenylamine is toxic and the indicator solution is strongly acidic. The dispensing bottle should have toxic and corrosive warning labels. While ferrocyanide and ferricyanide ions have fairly low toxicities, they can decompose in the presence of heat or strong acids to produce hydrogen cyanide. Precautions should be taken so that waste containing ferrocyanide and ferricyanide does not come into contact with strongly acidic solutions. W
Supplemental Material
A student handout providing background information, problems, instructions for performing the experiment, and a data sheet is available in this issue of JCE Online. Literature Cited 1. Murov, S. Experiments in General Chemistry, 3rd ed. (lab manual for Umland/Bellama’s General Chemistry); BrooksCole: Monterey, CA, 1999; pp 119–126; parts of that experiment are published here with the courteous permission of Thomson Learning. 2. Ginsburg, L. In Comprehensive Analytical Chemistry, Vol. Ic, Classical Analysis: Gravimetric and Titrimetric Determination of the Elements; Wilson, C. L.; Wilson, D. W., Eds; Elsevier: New York, 1962; pp 398–399. Kolthoff, I. M.; Sandell, E. B. Textbook of Quantitative Inorganic Analysis; Macmillan: New York, 1952; pp 549–555.
JChemEd.chem.wisc.edu • Vol. 78 No. 10 October 2001 • Journal of Chemical Education
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