infight Demonstration of Ionic Dissociation in Aqueous Solution Damon Dlemente Trinily School 101 West 91stSbeet New York. NY 10024
Introductory chemistry courses always include a discussion of the dissociation of ionic salts in aqueous solution. Such solutions contain independent, mobile cations and anions ~~-and are exoected to be eood conductors of electricitv. understandab];, therefore, many teachers illustrate the faEt of dissociation with a conductivitv demonstration. Unfortunately, however, the conductivitybf ionic solutions is a rather comnlex behavior. Electrons move through the wires, passive cations and anions move in opposite directions i n t h e solution, and redox reactions take place a t the electrode surfaces. In many ways a conductivity demonstration will raise more ouestions about other topics than it will answer about dissociation. So here is a simple, visual demonstration that readily convinces students that many ionic solids exist in adifferent form in solution from that which they exhihit in the solid ~
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Choose from the stockroom shelves a series of salts containing a common transition-metal ion. Given what is available tome, I usually select copper(I1) salts. I have found this set of copper(I1) salts to be ideal: the chloride dihydrate, the sulfate pentahydrate, the anhydrous sulfate, the acetate dihvdrate. the carbonate, the hydrated nitrate, and the anhydrous homide. Begin the demonstration by displaying samples of these solids to the class and asking for comments. Students immediately notice that most of these salts are some shade of blue or p i e n . The two exceptions are the anhydrous sulfate (white, though usually faintly tinged with blue) and the bromide (very dark, nearly black). When asked to account for this observation, a concensus is quickly reached. The similarity in the colors is attributed to the presence of the common ion copper(I1). But if coooer ion causes the color.. why. are the salts not identical sh'ades of blue? This is a more subtle inquiry, but the conclusion is alwavs reached that since the copper(I1) is in a different environment of anions and water mhiecules in each compound, i t is not surprising that the colors vary. Is i t possible to show that the bluish colors cannot be due to the anions? Simply display sodium or potassium salts containing the anions and note that they are all white solids and give colorless solutions. Now dissolve all the copper(I1) salts in water, and ask the class for comments. The first observation is that the carhonate is insoluble and so has to be dismissed from the rest of
950
Journal of Chemlcal Education
Wichita State University Wichita. KS 67208
the demonstration. The second and more critical observation is that the soluble salts eive solutions of identical color. (To make sure the colors are identical, arrange solute and solvent amounts so that all the solutions are about 0.1 M. This ensures equal color density and avoids complications due to ion pairing. Of course, too, the solutions should be made up in identical containers.) Why are the colors identical in thesolutions and not in the solids? The class concludes. thoueh this mav reauire a little suggestive guidance, that in each solution the copper(I1) ion is in the same environment, surrounded hy water and free of the anion. The demonstration can end a t this point, or some interesting structural details can be gone into. Copper(I1) ion in solution is described1 simply as square-planar Cu(HzO)a2+ or, with more precision, as C U ( H ~ O )with ~ ~ +a square plane of four waters close in and two longer axial bonds. In either case, the blue color of the solution is attributed to a copper(I1) ion with four to sixcovalent hondsto oxygen atoms in water molecules. The solution color is almost identical to the color of crystals of copper(I1) sulfate pentahydrate. And no wonder, for in that salt1 the copper is bonded to four oxygens from water molecules in a plane, and to two axial oxygens from sulfate ions. In the anhvdrous sulfate. this structure is disrupted and takes the colo; with it. Cwstals of hvdrated conner(I1) nitrate and of the hvdrated acetate arebf similar-but deeper color than thei; solutions. This suggests that these salts too have crystal structures in which the copper is bonded to many oxygens. Sure enough, in the case of the acetate, the solid is dimeric, with each copper having four bonds to acetate oxygens and one to a water m ~ l e c u l e . ~ The most anomalous crystal colors in this demonstration are the green of copper(I1) chloride dihydrate and the near hlack of the anhvdious bromide. a n d these have comer bonded to atomsLther than oxygen. The anhydrousA6romide3 has flat chains of coppers bonded to four bridging bromides:
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In the chloride dihydrate? each copper is bonded to four bridging chloride ions and two axial water molecules. I Cotton. F.: and Wllkinson. G. Advanced horganlc Chemistry, 2nd ed.; Intersclence: New York, 1966. Kochl, J.; Subramlan, R. Inorg. Chem. 1965, 4, 1527. Helrnholz, L. J. Am. Chem. Soc. 1947, 69. 886. 'Harker, D. Z.Krlst. 1936. 93. 136.
