In the Classroom
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tested demonstrations
George Gilbert Denison University Granville, OH 43023
A Demonstration of Crystal-Field Effects in Octahedral Complexes Submitted by:
Edward Koubek and Mark Elert Chemistry Department, U. S. Naval Academy, Annapolis, MD 21402
Checked by:
George Gilbert Denison University, Granville, OH 43023
A discussion of crystal field theory is usually included in general chemistry texts. This is likely to be one of only two places in the text—the other is the description of the hydrogen atom—where the important concept of light absorption by atoms and molecules is presented. The simple demonstration described here can perhaps enhance the presentation of crystal field splitting and magnetic properties in octahedral complexes.
Figure 1. Crystals of FeSO4?7H 2O are paramagnetic and adhere to the poles of a magnet.
Figure 2. Crystals of K4Fe(CN)6 are diamagnetic and do not adhere to the poles of a magnet.
After a discussion of octahedral crystal field splitting, the spectrochemical series, and high- and low-spin electronic configurations, students are shown two compounds containing Fe2+. One is pale green and the other is pale yellow. They are told that one is FeSO4?7H2O, which contains the Fe(H2O)62+ complex, and the other is K4Fe(CN)6. They are then asked to predict which compound is which, and also to predict which is more likely to be paramagnetic. After some discussion the instructor can place a strong magnet on an overhead projector with a petri dish underneath the poles. When the green crystals are sprinkled near the poles of the magnet, they adhere to the poles but can easily be brushed off and collected in the petri dish. The yellow crystals, on the other hand, have no attraction to the magnet. (See Figures 1 and 2.) Clearly, then, the green compound is paramagnetic and the yellow compound is diamagnetic. The weakness of the attraction in the paramagnetic case should be emphasized; this is an excellent opportunity to point out the difference between paramagnetic and ferromagnetic behavior. The differences in color and magnetic behavior between the two compounds can readily be explained on the basis of the crystal field model. Both solids contain octahedral complexes of Fe2+, which has a d6 configuration. In Fe(CN)64–, the strong CN– ligand gives rise to a large crystal field splitting ∆, and all six d electrons remain paired in the three lower t2g orbitals, producing a diamagnetic complex. Absorption occurs at the blue end of the visible spectrum, so the complex appears yellow. H2O appears lower in the spectrochemical series, and therefore it is not surprising that Fe(H2O)62+ is a weak-field, high-spin complex with four unpaired electrons. Absorption in the red, or low-energy, end of the spectrum causes the hexaquo complex to appear green. The crystal field splittings for Fe(CN) 64– and Fe(H2O)62+ are known to be 33,000 cm –1 and 10,400 cm–1, respectively (1). The centers of the d–d absorption bands therefore actually lie outside the visible part of the spectrum in both cases. In Fe(H2O)62+, the broad d–d absorption band clearly extends into the red part of the visible spectrum to give the observed green color. The situation for Fe(CN)64– is somewhat less clear. The d–d absorption is obscured by a strong charge transfer band, which begins to grow in below about 450 nm, and it is probably the latter that is responsible for the bulk of the visible light absorption. Nevertheless, the yellow color of the K4Fe(CN)6 clearly demonstrates that the crystal field splitting in this material is at or beyond the blue end of the visible spectrum. This demonstration also makes a good “explain the demo” test question for an upper-level inorganic course. Literature Cited 1. Cotton, F. A.; Wilkinson, G.; Gaus, P. L. Basic Inorganic Chemistry, 3rd ed.; Wiley: New York, 1995; p 516.
Vol. 73 No. 10 October 1996 • Journal of Chemical Education
947
