Metal-Ligand Complexes: Ni+2/en

section V-A

last edited: 9/3/01 (Web edition)

Spectrophotometry / Background

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).

Integrating yields the common form of Beer's law The term -log10[ I/I° ] is called A, the Absorbance and the constant 2.303k is renamed e, the molar extinction coefficient (Greek epsilon.) By convention the units of concentration are moles/liter and the path length, l, is in cm. The Absorbance has no units but is often written as a number followed by the letter A or the symbol AU (absorbance units.) The most common application is the determination of solution concentration by measuring the Absorbance. The molar extinction coefficient is either found in the literature or it is determined from measurements of Absorbance on solutions of known concentration. The path length, l, is generally the width of a cuvette and most cuvettes are a standard 1.00 cm.

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.


In this case it makes more sense to submit separate reports for Part I and Part II. They both refer to the spectrum of inorganic complexes, but are otherwise unrelated. The report for part I is typically 1-2 pages in length. Part II is more involved and will require mastery of spreadsheets and finding a way to present spreadsheet results.

Experimental: Part I


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.)

Prepare the solutions listed and measure the absorbance at 480 nm. Measure quickly and at a consistent time (say 1 or 2 minutes after mixing.) Caution-- make one solution and measure it before working on the next mixture. For the first solution, measure the absorbance several times over an hour or two to see if the reading is constant.


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 starting solutions are: The mixtures can be made by putting Ni2+ solution into a buret. For example, delivering 3.00 ml of Ni3+to a 10.00 m; volumetric flask and diluting to volume with en solution will produce a solution with XNi=0.30. Ten ml. is an adequate sample size for these determinations. The solutions will keep. They can be prepared in one period and measured in the next period.

the mixtures should have