In the Classroom edited by
Tested Demonstrations
Ed Vitz Kutztown University Kutztown, PA 19530
Purple or Colorless—Which Way Up? An Entertaining Solubility Demonstration submitted by:
Trevor M. Kitson Institute of Fundamental Sciences (Chemistry), Massey University, Palmerston North, New Zealand;
[email protected] checked by:
Bruce J. Heyen Department of Chemistry, Tabor College, Hillsboro, KS 67063
This demonstration can be quickly and easily prepared for use in a lecture or classroom. It has relevance to several key chemical concepts including polarity, intermolecular forces, solubility, and spectrophotometry. It is visually striking and with a little ‘showman’s patter’ it can be quite amusing. Several other demonstrations that involve solutions of iodine have been reported in this Journal (1–4); however, the special and valuable characteristic of the present contribution is in its humorous presentation. The entertainment factor in teaching science can be vitally important in building enthusiasm and aiding understanding. This demonstration amuses—and bemuses—students and for a number of years I have found it to be popular and well-received. Experiment Prepare four bottles containing the following and place them in a cardboard box: •
Bottle A: Hexane
•
Bottle B: An aqueous solution of potassium permanganate of sufficient concentration to give a deep color for visual impact. This solution should be acidified with a little dilute sulfuric acid otherwise it will slowly turn brown as the following reaction occurs: 4MnO4− + 2H2O → 4MnO2 + 3O2 + 4OH−
•
Bottle C: Water
•
Bottle D: A solution of iodine in hexane. The concentration should be carefully adjusted until this solution appears to the naked eye to be exactly the same color and intensity as the solution in bottle B.
Tell the students you are going to demonstrate the immiscibility of water and hexane and to make the interface more obvious you have colored the aqueous solution with potassium permanganate. Talk briefly about the solubility of many ionic compounds in water, a highly polar solvent, and their insolubility in nonpolar hexane. Pour samples from bottles A and B into a container (a boiling tube or, preferably on a large scale, a 500-mL measuring cylinder). Two layers result, a colorless layer above a purple layer. Mention that the laboratory technician whom you asked to prepare the solutions has apparently provided two samples of each, so you may as well use them. Repeat the demonstration, surreptitiously taking bottles C and D out of the box 892
instead of bottles A and B. Again two layers result, but this time the colorless one is below the purple one. Scratch your head in pretended mystification. Why didn’t I quit while I was ahead? Why do things always go wrong for me? What has the technician done differently? Class Discussion Encourage the students to think along with you as to what might be the explanation. With a bit of luck (or with prompting from you) someone may suggest that the second demonstration used a solvent denser than water, such as dichloro- or trichloromethane. (Hexane, of course, is less dense than water.) Tell them this is an excellent idea and can be easily tested by adding more water, which should simply dilute the purple layer. Add some water, preferably directly from the tap so that you can say this is surely genuine water, there is no deception involved here! The purple layer is unaffected, the lower colorless layer increases in volume. So the proposed explanation is wrong; the lower layer is definitely water. As a control experiment, add extra water to the first mixture too, and observe the expected dilution of the permanganate solution. In a similar way, show them that the expected result happens if more authentic hexane is added to the two mixtures. Yes, in both cases, the upper layer grows in size; it is hexane. So how on earth have we got the color purple to prefer water in one case and hexane in the other? Students’ suggestions may be quite imaginative here. Finally ‘discover’ a detached label at the bottom of the box, saying “Iodine in hexane.” Aha, that’s it! In the second case the color is not due to potassium permanganate at all, but to iodine. Talk about the lack of polarity of the iodine molecule and how its solubility in hexane illustrates the ‘like dissolves like’ rule of thumb. In case students are confused by their experience of the brown aqueous solutions of ‘iodine’ that they may have met in iodometry experiments, remind them that these solutions contain mainly the triiodide ion rather than iodine itself. Visible Spectra Show the students overhead transparencies of the visible spectra of the colored solutions you have used in this
Journal of Chemical Education • Vol. 80 No. 8 August 2003 • JChemEd.chem.wisc.edu
In the Classroom
Additional Insights For advanced students, you may like to describe how potassium permanganate can indeed be dissolved in nonpolar solvents in the presence of the crown ether, 18-crown-6, to sequester the potassium ion. This then makes a powerful reagent for the one-phase oxidation of alkenes, for example (5). It might also be mentioned that iodine dissolves appreciably in pure water giving a yellow-brown solution. Lewis bases such as water, alcohols, and amines tend to donate electron density into the iodine molecule’s LUMO. The resulting charge-transfer complexes absorb light at a shorter wavelength than the unperturbed free iodine, or iodine in a non-donating solvent such as hexane (6). (In the present demonstration iodine does not detectably partition from hexane into the water phase.) Interestingly, the color of the permanganate ion is also due to charge-transfer transitions, in this case between the oxygen (formal oxidation state ᎑2) and manganese (oxidation state +7); no d–d transitions, which explain the color of many transition metal complexes, can occur here. It is noteworthy that other corresponding oxyanions of group 7 of the periodic table (TcO4− and ReO4−) are colorless as the energy gaps of their charge-transfer transitions lie in the UV region (7). Hazards Figure 1. Visible spectra of (A) aqueous potassium permanganate and (B) iodine in hexane.
demonstration, again if you like ‘discovering’ that the technician has conveniently left these at the bottom of the box. Point out that although the absorbance and λmax values are essentially the same, the spectrophotometer can easily pick out the subtle differences in the spectra that are undetectable by eye (Figure 1). This might also be a good opportunity to demonstrate how much more precise a spectrophotometer is than the human eye at measuring the intensity of color (and hence concentration). For instance, you could show the students a previously prepared series of dilutions of potassium permanganate and ask them to estimate the factors by which the concentrations differ; then show them the corresponding spectra with the peak absorbance values measured accurately. (This demonstration is more striking with solutions of, say, the p-nitrophenoxide ion, as the eye is particularly bad at quantitatively estimating the color yellow.)
Suitable precautions should be taken with solid potassium permanganate and iodine, which are both oxidizing agents. After use, hexane and hexane兾iodine should be stored in a solvent waste bottle for later disposal and not poured down the sink. No sources of ignition should be nearby when using hexane. The demonstrator should wear the appropriate safety goggles. Literature Cited 1. Cliche, J.-M.; Labbe, B. J. Chem. Educ. 1988, 65, 813. 2. Nordstrom, B. H.; Lothrop, K. H. J. Chem Educ. 1984, 61, 1009. 3. Smith, W. L. J. Chem. Educ. 1977, 54, 228. 4. Summerlin, L. R. J. Chem. Educ. 1964, 41, A883. 5. McMurry, J. Organic Chemistry, 5th ed.; Brooks/Cole: Pacific Grove, CA, 1999; p 725. 6. Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. Advanced Inorganic Chemistry, 6th ed.; John Wiley & Sons, Inc.: New York, 1999; p 551–553. 7. Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth/Heinemann: Oxford, England, 1997; p 1050.
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