Absorption Spectroscopy

Absorption Spectroscopy is the most common tool in analytical chemistry.  Measurements are quick, accurate, inexpensive, and usually take little time to prepare.  However, although spectrophotometry is based in color, it is almost always taught with instruments that present a sterile black-and-white view of a conceptually colorful and interesting science.

Visual Absorption Spectroscopy

Students can use MicroLab’s Visual Spectrophotometer and our FASTspecTM scanning spectrometer to bridge the gap between visual and conceptual observations to quantitative measurements of absorbance spectra.

AS1Ask a freshman what is going on in this vial and they will say, “It’s giving off green light!”  The key to understanding absorbance spectroscopy is knowledge that color is not “given off”, but is what remains when some wavelengths are absorbed from a white spectrum.  Students more easily think in terms of percent transmittance—the colors that are passed through a sample.  But spectrophotometry works on absorbance—colors that are taken out.  Some quick experiments with our Visual Spectrometer and the FASTspecTM scanning spectrometer will make this point well.

Try holding a colored sample in front of “blank” spectra displayed on your computer screen below.  Then slowly move the sample across the spectrum.  Were you able to view the absorption bands?  Wasn’t that fun?

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If you place a vial of colored liquid in the Visual Spectrometer and look at a white light source (the sky, or an incandescent white light bulb), you can view the absorption bands together with a continuous white light reference spectrum.  Students learn that the food dye sample looks green because it absorbs blue/violet and red light.  These absorption bands are real and highly visual.

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Our new FASTspecTM scanning spectrophotometer software creates this display with every scan, providing students with a visual representation corresponding to the quantitative absorption data collected.  A photographic image of the “blank”, or white light reference, is at the top with a copy of the sample spectrum below.  This sample spectrum is created in the software with a black overlay.  The density of the black overlay is proportional to the absorbance at each wavelength.  Students can relate the transmission minima or absorption maxima on the graph to the actual absorption of light by the sample.

Quantitative Absorption Spectroscopy

From absorption, we move to Beer’s Law.  Beer’s Law is the workhorse of analytical chemistry.  It is a mathematical statement of the relationship between the absorbance of a solution (A) at a specific wavelength, the concentration of this solution (c), and the width of your sample (i.e., the path length the light will travel through your sample) (l).  The constant that relates these values is the molar absorptivity (e).  This coefficient tells us how well the molecule can interact and absorb light at a given wavelength.

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However, selection of a wavelength for Beer’s Law experiments is counter-intuitive or sometimes not even considered by students.  They find it hard to understand, again because it depends on the light absorbed, not on what light comes through the colored sample.  Conventional wisdom is to choose the wavelength that corresponds to the color of the sample.

We’ve taken our green sample above and quantitatively diluted it several times to obtain four samples and measured their absorbance.  Below are the resulting straight-line fits.  The difference between these graphs is the choice of analytical wavelength. In the first graph, we have selected 383 nm (violet) as our analytical wavelength.  Let’s compare this with our second graph where 502 nm (green) was selected.  You can see the change in absorbance at green is less than the change observed at violet, meaning green has lower sensitivity.  This is counter-intuitive, but expected because the absorbance of the sample is the lowest in the green region.  While both wavelengths give excellent fits, the difference in absorbance values are greater for violet, meaning you would be able to more accurately determine small concentration differences.

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The red (635 nm) measurement illustrates deviations from Beer’s Law at high absorbance.  When absorbance is high, very small amounts of stray light become significant and lower the observed absorbance at the detector.  A non-linear curve fit works for concentration, through the errors in measuring an unknown become much greater as the curve flattens out.

The FASTspecTM sample holder can accommodate vials of three diameters.  Students can compensate for more concentrated samples by choosing a vial with lesser path length or use them to explore the path length dependence of Beer’s Law.

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