03. Spectroscopy of Atoms and Molecules & Beer's Law

Figure 1: Visible Colors of the Electromagnetic Spectrum

The interaction of light with objects

One of the most obvious properties of any substance is its color. The green color of leaves is due to chlorophyll absorption, the orange of a carrot or tomato arises from carotenes, and the red of blood results from hemoglobin. For solid objects, the color is characteristic of the spectrum (in the visible region) of light scattered by the substance when white light (or sunlight) shines on it (light that is not absorbed/unabsorbed is scattered).

White light contains all of the visible colors in the electromagnetic spectrum. We can use a color wheel to help us understand that the color an objects absorbs (which we will not see) is complimentary to the color that we see.

Figure 2: The Color Wheel. Complementary colors are opposite each other.

To investigate what colors are absorbed or not absorbed by an object, we will use Gummi Bears! The Gummi Bears come in different colors (ex: red, orange, green, yellow, blue and white). You will use three different laser pointers (red, green, and purple or blue) to shine a monochromatic light (light with a single color and a single energy/wavelength) on the Gummi Bears to detect how each colored Gummi Bear absorbs, reflects, or transmits a specific color of light.

Absorption Spectra in Solution

We use an absorption spectrum to quantitatively measure the extent to which light is absorbed by a colored sample. For example, Chlorophyll is green because it transmits in the green region of the electromagnetic spectrum (550 nm), and it absorbs the complimentary colors in the blue (435 nm) and red (660 nm) regions of the spectrum, as seen in Figure 2.

Figure 3: The absorbance spectrum of chlorophyll.

Using a spectrophotometer to measure absorption spectra

In order to measure an absorption spectrum, a spectrophotometer, also known as a spectrometer, is used. A simple version of this instrument is shown in Figure 3. The light from a lamp is split into its component wavelengths by a prism or diffraction grating. This selected light is sent through our solution, and the detector then measures the amount of light that reached it. A computer compares how much light reached the detector when no sample was in place with how much light reached the detector when a sample was in place.

Figure 4: Schematic of a simple spectrophotometer

When you use a spectrophotometer, you must first calibrate the instrument (also known as "blanking" or "zeroing"). To do this, you make a "blank" solution that is identical to the solution you will measure but contains none of the analyte that you wish to measure. You put the blank solution in a plastic cuvette in the spectrophotometer and calibrate the instrument to the blank, setting the absorbance to zero when the blank is present. Thus, when you put your solution containing analyte into the spectrophotometer, it will only measure absorbance due to the analyte.

Beer's Law can be used to relate absorbance to concentration

The absorption spectrum can help us determine the amount of colored substance we have in a solution. We can make use of Beer’s Law, which states that absorbance of a solution is proportional to the concentration of absorbing atoms/molecules in the solution:

A = εlc

A = Absorbance (unitless)

ε = Molar extinction (or absorptivity) coefficient (cm^-1 M^-1)

l = Path length (cm)

c = Concentration (M)

The molar extinction coefficient is a constant that is different for each molecule. It also depends on which wavelength you measure the absorbance at (i.e., ε measured at 470 nm is different than ε measured at 475 nm, even for the same substance).

This equation should make intuitive sense: If the concentration of molecules increases, the absorbance should increase, because more molecules are interacting with light. Also, if the path length (l) increases, the absorbance increases for exactly the same reason. The cuvettes that we use in Chem 105 lab have a standard path length of 1.0 cm.

Figure 4: Beer's Law plot (aka-"standard curve")

In order to determine molar extinction coefficient, ε, you must make a calibration curve using a series of standard solutions (see Figure 4). A standard solution is a solution in which the concentration of the analyte is known, "c." The absorbances of several standard solutions are measured and absorbance vs. concentration is plotted. The plot is called a "standard curve", and should be a straight line with y-intercept virtually zero, whose equation y = mx + b corresponds to the Beer's Law equation A = εlc. (Why should the y-intercept be zero?) In these equations, y corresponds to A, x corresponds to c, l corresponds to the "path length of the cell in the spectrometer (1 cm, in our lab)and the slope corresponds to εl. Once the calibration curve is complete, the absorbance of the unknown sample is measured, and the value of the extinction coefficient (ε) is determined in Beer's Law using the slope of the standard curve: εl. The Beer's Law formula can then be used to determine the concentration of an unknown in solution, such as the amount of iron in a vitamin pill...next week's lab!

Using spectrophotometry to measure the amount of iron in a solution with known amounts of iron (this lab) and unknown amounts of iron (next week's lab)

In this lab, you will be using spectrophotometry to measure the amount of iron in various diluted samples creating known concentrations of iron. Next week's lab, we will use the data from this lab to determine the amount of iron in a vitamin pill.

Measuring iron spectrophotometrically: Making a colored solution

When iron is dissolved in water, the resulting solution is colorless. Spectrophotometry can only be used on a colored solution. In this lab, you will be performing a number of reactions, including iron ions reacting with 1,10- phenanthroline (aka-"phen") to form a colored ionic iron-phenanthroline complex to measure its absorbance:

Fe²⁺ + 3 Phen→ [Fe(Phen)₃]²⁺

Image from: http://water.me.vccs.edu/courses/env211/changes/ironreaction.gif

You will start with a colorless solution of ferrous ammonium sulfate (Fe(NH₄)₂(SO₄)₂ • 6H₂O), which dissociates in water to yield one Fe²⁺_(aq) ion for each ferrous ammonium sulfate molecule. You will perform a reaction to turn colorless Fe²⁺_(aq) into a colorful iron-phenanthroline complex which can be measured spectrophotometrically.

First, you need to ensure that all of the iron is in the Fe²⁺_(aq) form. Iron exists predominately in three forms: Fe_(s), Fe²⁺ and Fe³⁺. Acid has already been added to the ferrous ammonium sulfate solution to ensure that any Fe_(s) impurities are oxidized to Fe³⁺_(aq). Your first step will be to add a reducing agent, hydroxylamine hydrochloride (HONH₂•HCl), to ensure that any Fe³⁺_(aq) impurities are converted to Fe²⁺_(aq).

Next, you will add sodium acetate (NaOCOCH₃), a buffering reagent, to the ferrous ammonium sulfate solution to make sure that the solution has a pH that is appropriate for the next step: formation of the spectrally active iron-phenanthroline complex.

In the third step, you will react your Fe²⁺ solution with 1,10-phenanthroline (C₁₂N₂H₈). One Fe²⁺ ion forms a complex with three 1,10-phenanthroline molecules as seen here:

Fe²⁺ + 3 Phen→ [Fe(Phen)₃]²⁺

Fe²⁺ ions are the "limiting reagent" because we are adding excess Phenanthroline to our solutions. All of the Fe²⁺ ions are used up in the reaction and the amount of products formed are limited by the number of Fe²⁺ ions initially available.

In solution, the iron-phenanthroline complex is a bright orange color. You will measure the absorbance of your standard solutions. First, you will determine the wavelength of maximum absorbance, λ_max. This will be the wavelength that you will take all of your absorbance measurements at. Then, you will take the absorbance of multiple solutions of known concentrations. Plotting this data as A vs. conc. (M) will result in a linear fit with slope εl, according to Beer’s Law.

Making solutions: Know your glassware

In order to make our standard curve, we need to know the concentration of our analyte (the iron-phenanthroline complex; see "Synopsis" below) precisely. This means that we need to know 1) the precise volume, and 2) the precise amount of analyte. Since iron is our limiting reactant for all our solutions, it's quantity is the only one that needs to be measured precisely. All the other reactants are in excess, which means that there will be plenty of them available to react even if our amounts are inexact, so we do not need to measure the amounts precisely.

If you need to know a volume precisely, you need to use volumetric glassware (see Figure 5). If you do not need to know a volume precisely, you can use a graduated cylinder!

Figure 5: Spectrophotometer, Part 2 Glassware, & Part 1-Gummi Bear station

Synopsis

There are two portions to this lab:

In Part I, you will use different colored Gummi Bears to hypothesize and then investigate what colored light each colored Gummi Bear best absorbs or does NOT absorb due to the color of the Gummi Bear.

In Part II, you will make a standard curve for an iron-phenanthroline complex (which you will use next week to determine the iron content in a vitamin pill.) The goal of this experiment is to determine the molar extinction coefficient ε for the iron-phenanthroline complex using the Beer's Law standard curve. In our next lab, we will use your standard curve from this lab to confirm or refute the amount of iron listed on the ingredients label for a vitamin pill.

Experiment

To print instructions, select the portion that you wish to print, right click and choose "print" to prevent printing the entire web page.

Materials and Equipment for Part 1:

• • Laser pointers (red, green, and blue/purple)

  • Gummi Bears

Materials and Equipment for Part 2:

• • 2% hydroxylamine hydrochloride, HONH₂• HCl

  • 10% sodium acetate, CH₃COO^- Na⁺

  • 5 x 10^-4 M solution of ferrous ammonium sulfate in 1% sulfuric acid (Record exact molarity to 4 sig figs)

  • 0.002 M 1,10-phenanthroline, C₁₂N₂H₈

  • Ocean Optics spectrophotometer

  • 50.00 mL volumetric flasks

  • 1.00-8.00 mL volumetric pipettes

  • Various sizes of graduated cylinders

Instructions

Students will work in pre-assigned pairs. Parts I and II may be done in any order.

Part I: Laser pointers and Gummi Bears

To study the interaction of light with objects, you will use Gummi Bears with various colors. The light sources will be laser pointers with red, green, and blue/purple lights. These lasers have a very narrow spectral range of wavelengths of light produced (e.g. the green laser is mostly photons with 532 nm).

In the table below, predict how the different colored laser pointers will interact with the Gummi Bears listed. Also, use this table to record your observations when doing the experiment.

OPTIONAL: Please feel free to expand the data table below to include any other colored Gummi bears that are present in your lab.

Warning: You will be using high powered laser pointers and if possible, please use the fume hood for this experiment or a place where you can point the laser pointers away from any students in your lab. No matter what, the following safety rules must be observed:

• Engage the laser pointers ONLY when you are making an observation and if it is properly positioned.

• Make sure there is no reflective surface (e.g. glass) in the path of the laser beam at any distance.

• Make sure there is no person in the path of the beam.

• Engage the laser beam only for a short period of time, until you can make an observation.

Follow these instructions to complete the data table:

1. Point the Gummi Bear's head towards the back of the hood

2. Position one of the lasers at the feet of the first Gummi Bear so that the beam will propagate through the length of his body and out his head.

3. Turn on the laser pointer by pushing the button.

4. Observe the interaction of the laser beam as it goes through the Gummi Bear:

   • If no absorption takes place, the laser beam will scatter through the body of the Gummi Bear and the Gummi Bear will light up.

   • If absorption takes place, the Gummi Bear should appear "unlit" and/or much less illuminated then those Gummi Bears that 'light up.'

Please note all your observations and conclusions in the data table above. Please feel free to create a third column for any data you want to include about the purple/blue laser pointer.

References

¹Part of this lab is a modified version of "Atomic Spectroscopy" In Experimental Chemistry by James Hall, Houghton Mifflin Co., Boston 1997.

²Wikipedia, Prism (optics) 2/13/08.