In the Classroom Tested Demonstrations
“Mud” + “Blood”— A Very Colorful Demonstration submitted by:
Gordon Hambly John Abbott College, P.O. Box 2000, Ste. Anne de Bellevue, PQ H9X 3L9, Canada
checked by:
James Ealy The Peddie School, Hightstown, NJ 08520
Abstract A red solution of phenolphthalein indicator in base is added to a muddy manganese dioxide solution, and a clear solution is obtained almost immediately. This dramatic demonstration involves both redox and acid–base chemistry and can serve as a review of these subjects. Keywords Demonstrations Redox Reactions Supplementary Materials No supplementary material available.
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Journal of Chemical Education • Vol. 75 No. 1 January 1998 • JChemEd.chem.wisc.edu
In the Classroom Tested Demonstrations
“Mud” + “Blood”— A Very Colorful Demonstration submitted by:
Gordon Hambly John Abbott College, P.O. Box 2000, Ste. Anne de Bellevue, PQ H9X 3L9, Canada
checked by:
James Ealy The Peddie School, Hightstown, NJ 08520
A central tenet of Christianity, and Judaism in old-testament times, is that the forgiveness of sin involves the application of blood (Leviticus 16:11–16; Hebrews 9:22; 1 John 1:7). I was asked to design a demonstration that would illustrate this principle. Specifically, a red solution (the blood) is poured into a dark mixture (sin) and a colorless solution must be obtained quickly. I’ll describe the demonstration first and then explain how I use it in my secular classroom. Experimental Procedure Goggles, an apron, and gloves should be worn for safety during the demonstration. The 2 M sulfuric acid used is relatively dilute (concentrated sulfuric acid is about 18 M). Into a 400-mL beaker are placed about 0.1 g of potassium permanganate (KMnO4) and about 0.1 g of hydrated manganous chloride (MnCl2?4H2O) or an equimolar amount of any water-soluble manganese(II) salt, along with 100 mL of water and 50 mL of 2 M sulfuric acid. The result is a muddy brown-purple mixture. In a 250-mL beaker are added 10 mL of 3% hydrogen peroxide, 100 mL of water, 20 mL of 1 M sodium hydroxide, and 10 drops of phenolphthalein indicator solution. The indicator in the base will appear deep pink, relatively close to blood red. The “blood” solution is now added to the muddy “sin” mixture. A clear solution is obtained almost immediately. The excess sulfuric acid can be neutralized with base (the phenolphthalein indicator is already present) before disposal. In my classroom in a secular college, I introduce the demonstration by briefly alluding to the religious back-
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ground noted in the first sentence of this report. I then perform the demonstration without further explanation. Next I ask my students two questions. First, given the partial chemical reaction below, what is the second product? (I get some very interesting responses, especially if redox has not been covered recently): MnO 4{ + H2O2 → Mn2+ + ? The permanganate ion is reduced while the peroxide is oxidized to oxygen gas. The balancing of this redox reaction can be reviewed. The permanganate ion is purple, whereas the brown color is due to manganese dioxide, which is also reduced to manganese(II) ion in the reaction. Some manganese dioxide is present as an impurity in the permanganate (1), but most of it is produced by reaction of the permanganate and manganous ions. The final solution is colorless because the peroxide is present in excess and the products of the reaction are colorless at the concentrations used. Second, I ask my students how I was able to color the peroxide solution reddish and then have this color disappear. Be prepared to give hints such as, “Have you ever seen anything change color in the lab?” With enough appropriate hints, someone will come up with the idea of the phenolphthalein acid–base indicator. The excess acid in the permanganate solution not only ensures that the phenolphthalein indicator will end up colorless, but also directs the reduction of the permanganate specifically to the colorless manganous ion rather than to brown manganese dioxide. Literature Cited 1. Kolthoff, I. M.; Sandell, E. B.; Meehan, E. J.; Bruckenstein, S. Qualitative Chemical Analysis, 4th ed.; Macmillan: London, 1969; p 817.
Journal of Chemical Education • Vol. 75 No. 1 January 1998 • JChemEd.chem.wisc.edu
