Naked Science Forum
Non Life Sciences => Chemistry => Topic started by: valeg96 on 14/03/2014 18:23:32
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So I've prepared a couple of experiments for a self-management day at school. I'll have a complete school lab and a teacher for help and surveillance. I thought i'd do the street light bottle (indigotine+glucose+Na/KOH), the blue bottle (methylene blue, glucose, NaOH) and a quick show of metal/acid reactions and the reactions of CuSO4 with ammonia, NaOH, HCl, Na2CO3, and KBr (was it? haven't tried that one yet). The issue is i have 4 HOURS, and aside explaining the basic rules in some interesting and funny way (no eating, no cooking etc) and showing what actually happens when you drop water on 98% sulfuric acid, that's pretty much it. I'm looking for colorful, funny and unusual experiments, with a meaning and that can be explained (no random mixing of various things like the elephant toothpaste) and that may be joined to give the whole thing some sense. I thought of explaining how the basic reagents can be produced, maybe starting from an apparently pointless reaction and going up like (not balanced and complete)
C6H8O7+Ca(OH)2->Ca3(C6H5O7)2 (precipitation of calcium citrate)
NaOH+CaCl2->Ca(OH)2 (production of calcium hydroxide)
Na+H2O->NaOH (production of sodium hydroxide, may try the electrolysis too)
CaCO3+HCl->CaCl2
Ca(OH)2+CO2->CaCO3
Ca+H2O->Ca(OH)2 (pointless loop)
This way I'd get to talk about acids/bases reactivity and get them to understand some REALLY GENERAL APPROXIMATE-ISH concepts.
Do you have any idea on how to complete this/carry this out? The key point is to keep them interested but not end up doing pointless things like these tv programmes that show random experiments to make people say "OOOH SCIENCEE LULZ"
Thanks a lot :3
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I recently had luck demonstrating catalysis of hydrogen peroxide decomposition using different catalysts. I added 5 mg of each catalyst to a 250 mL flask with 20 mL of hydrogen peroxide (~10% v/v in water).
FeCl3--active catalyst is a dark species, so when a speck of FeCl3 is added, the solution first turns dark, then begins bubbling. This reaction picks up speed and may eventually reach a boil (runaway reaction) so be prepared. Once the H2O2 is all consumed, the solution returns to pale yellow/clear.
NaI--again, there is a color change to dark brown. Steady fast bubbling, then the bubbling stops and the solution turns clear again when the peroxide is gone. (actually I think this one is the same as "elephant's toothpaste," except it has no detergent added)
MnO2--this catalyst is heterogeneous (it never dissolves) but the reaction is immediate and very fast (by increasing the amount of catalyst added or the concentration of the peroxide, this can decompose almost instantaneously, which makes for a great demo, but you have to be very careful!)
Freshly chopped potatoes (5g)--there are many biological enzymes that catalyze hydrogen peroxide decomposition. I found potatoes to be easiest to work with, but horseradish and liver (or blood) also work very well. All biological catalysts have to be fresh because the enzymes denature fairly quickly.
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I had the luxury of demonstrating a Iodine clock reaction recently to some new Gymnasium students and it seemed to be appreciated, I suggest searching it!
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Another nifty demo is the temperature-dependence of the color of elemental sulfur (sulphur).
Sulfur appears bright yellow at room temperature, but when chilled with liquid nitrogen or dry ice, it become white. As it warms up to rt, the color grows in slowly. Heating sulfur above rt leads to darkening to orange. At its melting point (115 °C) it is dark red (like blood), and it darkens to black on further heating. If the sulfur is allowed to cool before reaching this black stage, or allowed to cool very slowly from the black stage, it will return to yellow. Rapid cooling of the black sulfur yields "plastic sulfur" a stretchy brown-grey substance that very slowly reverts back to crumbly yellow sulfur.
Note: this demo is best done with about 1 g powdered sulfur in a large test tube. The sulfur might meed to be protected from air at the very high temperatures (or risk producing SO2, which smells awful and is fairly toxic), but DON'T seal the tube tightly before heating it (a balloon full of CO2 or N2 tightly sealed over the top of the test tube is perfect)
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Oh, and the reason the sulfur changes color has to do with vibrations in the molecule broadening the range of light frequencies that can be absorbed. Cold sulfur absorbs very strongly in the UV, but not in the visible. As the molecule warms up, this absorption broadens, extending into the visible range. When it absorbs only violet and blue light, the sulfur appears yellow. As it begins absorbing green and then yellow light, it shifts through orange to red, and finally it absorbs the whole visible range (though the peak is still centered in the UV region) and appears black.