COMPLEX IONS AND THE SPECTROCHEMICAL SERIES

PURPOSE

DESCRIPTION

TIME REQUIRED

MATERIALS

Chemicals

copper sulfate pentahydrate

nickel sulfate heptahydrate

cobalt sulfate heptahydrate

95% ethyl alcohol

possible sources of ligands:

hydrochloric acid *

ammonia*

sodium thiosulfate*

oxalic acid*

acetic acid*

tartaric acid*

EDTA

sodium bromide

ethylenediamine

dimethylglyoxime

Equipment

pH paper

small test tubes

eyedroppers

HAZARDS

MODIFICATIONS/SUBSTITUTIONS

1. Copper sulfate pentahydrate is available from a garden supply store as root killer.
2. Possible substitutions for the ligands are: hydrochloric acid--muriatic acid available from a hardware store, ammonia solution--household ammonia available from a grocery store, sodium thiosulfate-- photo fixer from a camera/photo supply store, oxalic acid--ZUD from a grocery store, acetic acid--vinegar from a grocery store, and tartaric acid--cream of tartar from a grocery store.

PROCEDURE

1. Prepare the following cation test solutions: 1.0 M CuSO₄ (dissolve 250 g CuSO₄·5H₂O in sufficient distilled or deionized water to make 1.0 L of solution), 0.10 M NiSO₄ (dissolve 28.1 g of NiSO₄·7H₂O in sufficient distilled or deionized water to make 1.0 L of solution), and 0.10 M CoSO₄ (dissolve 28.1 g CoSO₄·7H₂O in sufficient water to make 1.0 L of solution).
2. Prepare saturated solutions of tartaric acid and of oxalic acid.
3. Adjust the pH of the solutions prepared in step 2 so that they are neutral to pH paper by adding small amounts of ammonia solution.
4. Prepare a 1% solution of dimethylglyoxime in ethyl alcohol solution.
5. Measure l ml of the copper(II) solution into 11 small test tubes. Repeat with the cobalt(II) and nickel(II) solutions.
6. Set aside a test tube of each metallic cation as a control.
7. Add several drops of each ligand solution or a small scoop of a solid ligand to each of the remaining 10 test tubes containing the copper(II) ion. Mix the contents of each tube thoroughly after each addition. Repeat for the cobalt and nickel solutions. Observe carefully as the ligands are added. Continue to add ligand until a noticeable color change occurs. If no color change has been observed after an excess of ligand (approx. 3 mL) has been added, record no visible change on the data chart. Always compare the sample with the control. Note: Some of the complexes may have several intermediate states in which the color may vary; therefore, be sure to add an excess of ligand after the initial color change.
8. Line up the resulting complexes for each cation in order of the spectrum (ROYGBIV).
9. Record observed color changes in a data table similar to the one following.
   
   

   		Copper+2		Cobalt+2		Nickel+2
   Ligand Added		Initial Color	New Color	Initial Color	New Color	Initial Color	New Color

   
   
10. 10 Determine the order of the spectrochemical series using information in the discussion below.

    DISPOSAL

    Solutions may be flushed down the drain with ample water.

    DISCUSSION

    In water solutions, transition metal ions are usually not found as single ions but as complex ions in which the metal ion is bonded to 2, 4, or 6 water molecules with coordinate covalent bonds. If the water solution is mixed with another species which can donate electrons to the metal, a new complex ion with the new species may be formed. The various species which can complex with metal ions are referred to as ligands. Often the color of the new complex ion is different from the original color of the metal ion in water.

    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.
    
    Source: Shakhashiri

    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.

    REFERENCES

    Beren and Brady, Laboratory Manual for General Chemistry Principles and Structure, Wiley and Sons, New York, 1986. A similar experiment is described.

    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.

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    Submitted by Eva Lou Apel, Larry Ferguson, Glenda Marshman, Regina Monks, Sam Sakurada, and Joe Trebella
    
    

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