|Ligand Added||Initial Color||New Color||Initial Color||New Color||Initial Color||New Color|
The color of the complex arises from the crystal field splitting of the five degenerate d orbitals into 2 energy levels by certain liqands.
The energy gap, labeled E, between the d orbitals is on the same order of magnitude as the energy of a photon of visible light. The actual magnitude of E for a given complex and therefore the color of the complex is determined by the following: which ligand is used; which metal is used and the oxidation state of the metal. Thus, the color of the complex ion can be used to determine the wavelength of light absorbed by using the color rosette shown.
The colors are arranged in a circle so that absorption of one color allows the color which is opposite on the rosette to become the visible color observed. For example, if an object appears to have a yellow color it is absorbing deep blue light; a red object is absorbing green light; etc. The wavelengths of light corresponding to their colors are also shown on the rosette. Of the wavelengths shown, 400 nm corresponds to the greatest energy and 720 nm corresponds to the smallest energy.
By determining the color absorbed by a series of ligands and the corresponding energy involved, the ligands can be ranked according to strength based on the relative amounts of energy absorbed (E). Such a ranking is called a spectrochemical series. The following version of the series shows the ligands in order of increasing E.
iodide ion < bromide ion < chloride ion < thiocyanate ion < fluoride ion < hydroxide ion < acetate ion < oxalate ion < water < ammonia < ethyendiamine < sulfite ion < nitrite ion < cyanide ion
The results obtained in this experiment for the ligands used with the copper(II), nickel(II), and cobalt(II) ions generally agree with the series as given above. Consumer chemicals, when available, were used as ligands. The weak acid solutions must be neutralized or made slightly basic with ammonium hydroxide before use as a ligand.
Possible questions and/or extensions include having students write formulas for the complexes formed, give names of the complexes formed, and correlate the strengths of the ligands as shown by the spectrochemical series results with the instability constants of the complex ions. Note, the strength of the ligand is not the only determining factor for instability constants; however, bidentate and tridentate ligands are generally bonded to the metal more tightly than monodentate liqands.
Brady and Humiston, General Chemistry, Principles and Structure, Wiley and Sons, New York, 1986. Crystal field theory and complex ions are discussed.
Brown and LeMay, Chemistry: The Central Science, Prentice-Hall, Englewood Cliffs, NJ, 1981, p. 723. Crystal field theory and complex ions are discussed.
Shakhashiri, B.Z., Chemical Demonstrations, Volume 1, University of Wisconsin Press, Madison, 1983, p. 260. This work describes the theory of the color of complexes.
Submitted by Eva Lou Apel, Larry Ferguson, Glenda Marshman, Regina Monks, Sam Sakurada, and Joe Trebella