Metal-Ligand Complexes: Ni+2/en
last edited: 9/3/01 (Web edition)
The Lambert-Beer Law states that the absorption of light in a thin layer of sample is proportional to the intensity of the light (I), the concentration of the species (c), and the path length (dx).
Most solutions satisfy Beer's law with a high degree of accuracy. The most common failure occurs when the species being observed participates in an equilibrium or is unstable. The failure is not really in Beer's law, but in the fact that the "c" we use is not the actual concentration in solution. If Beer's law isn't satisfied exactly, a calibration curve (Absorbance vs. concentration) still permits determining concentration.
Absorbance is a log unit. Absorbances of 0.1, 1.0, 2.0, and 3.0 mean that the sample absorbs 10%, 90%, 99% and 99.9% of the incident light, respectively. At high absorbances (above 1.5) the detector must accurately measure very low levels of light, and detector noise and stray light become important errors. At low absorbances (below 0.1) the detector must accurately report very minute changes in the light intensity.
As a practical point, absorbance readings between 0.1 < Abs < 1.5 tend to be more accurate and one should try to use solutions within this range. Readings below 0.01 or above 2.5 generally have very little accuracy or validity.
SPECTROMETRY OF THE Fe+3 + SCN- COMPLEX
Most inorganic species do not absorb strongly and direct spectrophotometric methods are not very useful. Even worse, other colored species present in a sample can interfere. A common technique is to force the species of interest to react with a selected reagent to form a strongly colored species which can be measured. Ideally, the reagent is selective and only produces one colored complex. It is also important that the reaction go to completion so that all of the species of interest contributes to the final measurement.
We will examine the use of SCN- ion as a reagent for the measurement of iron. As we will see this is not an example of a good spectrophotometric reagent. The complex has a tendency to fade (an undesirable characteristic.) The equilibrium constant for the formation is relatively small, and under some conditions a modest fraction of the Fe+3 remains uncomplexed. In practice we can use the deviations from Beer's law to estimate the equilibrium constant for this complex. (The reagent 1,10 phenanthroline is a much better choice for an iron determination.)
Plot the Absorbance vs. [SCN-]. A simple minded application of Beer's Law would call for all readings to be the same since there is the same amount of iron and an excess of SCN-. In practice this reaction requires a large excess of SCN- before you can assume that all the iron has been complexed.
Extrapolate your Absorbance vs. concentration plot to estimate the absorbance in the presence of an excess of SCN-; from that value, estimate the extinction coefficient of the complex. For your solutions, determine the degree to which the iron is complexed and estimate the equilibrium constant. (These will not be very precise, but you should be able to get the correct magnitude for Keq. More precise calculations would require activity coefficients and other corrections.)
(One source, Sime, Physical Chemistry, gives a value of Keq=139.) II. The
Stoichiometry and Spectra of Complex Ions
You will determine the absorption spectra of the three complexes formed when Ni+2 reacts with ethylenediamine. (We will abbreviate ethylenediamine as en in the rest of these notes.) This will be complicated since most of the solutions will contain a mixture of two species.
You will also use the absorption data to make a Job's Law plot and demonstrate the formation of all three Ni2+/en complexes. Job's method consists of seeing how specific properties (like the spectra) change as we vary the ratio of the two reactants, holding the sum of the two species constant.
the mixtures should have
Use the HP8452a Diode Array spectrophotometer to record the spectrum of each solution (the mixtures, the Ni2+ stock, and the en stock.)
For each absorbance reading (each mixture at each wavelength) compute
The quantity Y would be zero if no reaction occurred (we assume the absorbance of EN is negligible.) Thus Y is a measure of the deviation from Beer's law at a given wavelength due to reactions and equilibrium. For each wavelength, prepare a graph of Y vs Xen. If a complex with the formula Ni(en)n2+ is formed, the Y plot (at a suitable wavelength) will show a maximum (or minimum.)
(The lab Web site or CD-ROM will contain additional information and examples on the Job's Method.)
Analysis of the Data-- Spectra of Ni(en)n
Again, this will involve a lot of calculations. However, you must begin by recognizing the underlying problem. Few, if any, of the solutions you analyzed will contain a single chemical species. The spectra are, in general, the spectra of mixtures. Your goal is to extract the spectra of the three Ni(en)n complexes.
There are two extreme cases where we can extract the spectrum of a single species. The obvious case is pure Ni2+ ion. Less obvious is the Ni(en)32+ species. This is the only Ni species present when Xen >0.8 ; the equilibrium constants are large enough that we can assume complete reaction with an excess of reagent.
You need to convert the Absorbance data for these solutions into the molar extinction coefficient at each wavelength.
The solutions in the 0.1<Xen<0.45 region are dominated by two species: Ni2+ and Ni(en)2+ . It is a good approximation to assume that the concentration of each is fixed by stoichiometry:
Compare your spectra with those in Gmelin, Handbuch der Anorganische Chemie. A copy of the relevant sections can be found in the lab (or on Web site and Class CD-ROM.)
The stability constants (data from Gmelin) for the three Ni2+ / en species are: