Absorption Spectroscopy

Introduction

Why does the color of a copper chloride solution appear blue? As the white light hits the paint, which colors does the solution absorb and which colors does it transmit? In this activity students will observe the basic principles of absorption spectroscopy based on absorbance and transmittance of visible light.

Concepts

• Spectroscopy
• Visible light spectrum
• Absorbance and transmittance

Background

The visible light spectrum (380 nm−750 nm) is the light we are able to see. This spectrum is often referred to as “ROY G BIV” as a mnemonic device for the order of colors it produces. Violet has the shortest wavelength (about 400 nm) and red has the longest wavelength (about 625–740 nm).

Many common chemical solutions can be used as filters to demonstrate the principles of absorption and transmittance of visible light in the electromagnetic spectrum. Copper(II) chloride, ammonium dichromate, iron(III) chloride, potassium permanganate are all different colors because they absorb different wavelengths of visible light.

In this demonstration, students will observe the principles of absorption spectroscopy using a variety of different colored solutions. This activity was originally designed using all chemical solutions as filters. In an effort to create a safer lab, deionized water and food coloring have been substituted for the more hazardous chemical solutions.

Materials

Safety Precautions

Copper(II) chloride solution is highly toxic by ingestion and inhalation. Potassium permanganate solution is a strong oxidizing agent. It is irritating to the skin and eyes and slightly toxic by ingestion. Wear chemical splash goggles and chemical-resistant gloves while preparing solutions. Wash hands thoroughly with soap and water before leaving the laboratory. Follow all laboratory safety guidelines. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding.Potassium permanganate may be disposed of according to Flinn Suggested Disposal Method #12a. Copper(II) chloride may be disposed of according to Flinn Disposal Method #26b.

Prelab Preparation

Solutions
The six solutions needed for this demonstration should be either transferred from their shipping containers into the micro-chemistry solution bottle or prepared fresh and placed in the corresponding microchemistry solution bottle. Fill each bottle as close to the top as possible without spilling. Any extraneous room will cause air bubbles to show when the solution is turned on its side and can interfere with the electromagnetic spectrum appearance. 

1. Using a permanent marker, label the six microchemistry bottles (on the upper side of the bottle) of potassium permanganate, copper(II) chloride, yellow solution, orange solution, erbium chloride and praseodymium chloride.
2. Transfer the potassium permanganate solution to the microchemistry bottle labeled potassium permanganate.
3. If the copper(II) chloride solution is clear, pour the solution into the appropriate microchemistry bottle. If a precipitate or residue has formed in the solution, filter the precipitate using a funnel and filter paper before transferring the solution to the microchemistry bottle.
4. Yellow solution is prepared by placing 200 mL of deionized water into a 250-mL beaker. Add one drop of yellow food dye and stir. Transfer into the appropriate microchemistry bottle. Pour leftover liquid in the sink and rinse the beaker.
5. Orange solution is prepared by placing 200 mL of deionized water into a 250-mL beaker. Add two drops of red food dye and stir. Transfer into the microchemistry bottle labeled “Orange Solution.” Pour the leftover liquid in the sink and rinse the beaker.
6. Transfer the erbium chloride solution into the microchemistry bottle labeled “Erbium Chloride Solution.”
7. Transfer the praseodymium chloride solution into the microchemistry bottle labeled “Praseodymium Chloride Solution.” 

Apparatus

1. Obtain two black sheets of construction paper.
2. Form a 2-cm wide slit in the center of the stage of the overhead projector with the two sheets of black construction paper. Position the slit on the stage so that the image of the slit projected onto the projection screen is vertical (refer to Figures 1 and 2).
3. Place the holographic diffraction grating film above the lens of the overhead projector (see Figure 1).
   {12979_Preparation_Figure_1_Overhead projector}
   Wear gloves when handling the diffraction grating. If the spectra are not projected to the left and right of the screen, rotate the diffraction grating 90 degrees (the alignment of the grating is important). Once two bright spectra (mirror images of each other) are displayed horizontally to the left and right of the film screen (Figure 2), the diffraction grating should be secured to the lens with tape. Tape only on the outside edges of the diffraction grating and make sure the diffraction grating is flat.
   {12979_Preparation_Figure_2_Projected spectra}
4. Adjust the focus of the overhead projector so that the image of the slit is in sharp focus on the screen. The two spectra should also come into focus.
5. Once the slit is aligned so the “selected” spectrum is in the proper location to be viewed by the students, the construction paper should be secured to the overhead projector with tape.
6. Several activities can be performed to demonstrate that white light is composed of all colors, and many other properties of light.

Procedure

1. Place the overhead projector 10–15 feet from the projection screen.
2. Turn on the overhead projector and aim the light at the projection screen. Eliminate as much light in the room as possible by turning off lights and closing window blinds or curtains. Note: The visibility of the spectrum will be best in a very dark room.
3. Position the overhead so one of the two vertical spectra is on the screen or flat wall. Rotate the overhead stand 15–20˚ to the right or left so the spectrum is centered.
4. Have students observe the appearance of the visible spectrum without any filters in place.
5. Ask students to sketch the observed spectrum on the Absorption Spectrum Worksheet using colored pencils.
6. Place the orange solution in the middle of the overhead projector stage across the 2 cm slit opening. The bottle should be placed on its side, cap pointing either left or right, not up and down, so that the side with the most surface area of the microchemistry bottle is touching the overhead.
7. Have students draw the color spectrum they observe in the appropriate box of the worksheet.
8. Discuss with students which colors are absorbed (blue and violet—hence the black spot where those colors used to be) and which colors are being transmitted (green, yellow, orange and red—the colors that are still visible).
9. Remove the orange solution and repeat steps 6–8 three more times using the copper(II) chloride, potassium permanganate, and yellow solutions.
10. Remove all bottles from the overhead projector screen. Hold up the praseodymium chloride solution. Have students note the color and predict what they expect to see based on the results obtained with the other solutions.
11. Initial observations will be different as all the colors in the spectrum are present, as well as several dark, fine lines. Closer examination will reveal two significant lines in the blue-violet region. Have students sketch the absorption lines on the worksheet.
12. Repeat steps 9 and 10 using the erbium chloride solution.

Student Worksheet PDF

12979_Student1.pdf

Teacher Tips

• This kit contains enough materials to carry out the demonstration multiple times.
• The demonstration can reasonably be completed in one 50-minute class period.
• As an extension to this activity allow students to make “artificial” purple (potassium permanganate) and blue (copper(II) chloride) solution and see if they yield the expected results when used as filters.
• To optimize spectrum viewing, the microchemistry bottles should be as full as possible. Once each solution has been transferred to the microchemistry bottle, fill any remaining space with distilled or deionized water using a Beral pipet.

Sample Data

Observations
The sketches should resemble the electromagnetic spectrum with the appropriate colors that are observed or transmitted shown in color. The regions (colors) of the spectrum that are absorbed will appear black or very dark. Following is a list of the colors that should appear in each box.

• Potassium permanganate (blue solution)—blue, indigo, violet, red and orange are observed (transmitted).
• Copper(II) chloride (red solution)—red, orange, yellow, green, blue and violet are observed.
• Orange solution—green, yellow, orange and red are observed.
• Yellow solution—red, orange, yellow, green and blue are observed.
• Praseodymium chloride— absorption line present between the orange and yellow, there are additional lines in the blue and violet.
• Erbium chloride—major absorption lines in the green with faint lines in the red and blue

Answers to Questions

1. Compare the relationship between the blue solution and orange solution in terms of their absorbance of visible light and give the definition of complementary colors.

   Blue and orange are complements of each other. Thus, the orange solution absorbs blue light and the blue solution absorbs orange light.

2. If a green solution had been placed on the overhead, predict which colors of light would be absorbed and which colors would be allowed to pass through.

   Red would be absorbed and green, yellow, orange and violet would be transmitted.

3. What do the spectra of the rare earth elements, erbium and praseodymium, demonstrate?

   They demonstrate fine line absorption spectra. The absorption lines appear to energy transitions involving electrons in the f orbitals.

4. Did erbium and praseodymium yield the results you initially predicted? Why or why not?

   The erbium did support the hypothesis because there were major absorption lines in the green, yellow green and faint lines in the red and blue. The praseodymium did not support the hypothesis that it should absorb bands in the red/orange region. There was a strong absorption line between the orange and yellow; there is also an absorption line in the blue and violet regions.

Discussion

White light is composed of wavelengths of light from the visible spectrum as well as light wavelengths that are invisible to our eyes (i.e., infrared and ultraviolet). The visible spectrum is often referred to as “ROY G BIV” after the colors of light that are produced when white light is transmitted through a prism: Red, Orange, Yellow, Green, Blue, Indigo and Violet.

The color of an object results from the reflection of light from that object. Grass appears green when exposed to white light because the “green-colored” light waves that compose the white light are reflected from the surface of the grass and the other wavelengths (e.g., “blue,” “red,” “yellow”) are absorbed. The green wavelengths of light reflect back to our eyes and interact with the cones in the retina of the eye. The cones are the color receptors of the eye. Our brain receives the signals sent from these cones and interprets that our eyes are seeing the color green. However, most materials do not reflect a pure single wavelength color and absorb all the other colors in the spectrum. Most materials reflect or transmit a combination of colors from the visible spectrum; this gives objects their own distinctive colors. Wavelengths and intensities of light generate an enormous variety of colors that we see.

The solutions used in this demonstration are examples of white light filters. Thus, a solution that appears yellow absorbs mostly violet light—most of the other colors in the visible spectrum are transmitted and thus observed in the projected spectra. See the Observations for a description of the relationship between the color of each solution and the colors that are absorbed. Spectroscopy is defined as the measurement of the amount or intensity of light or electromagnetic energy absorbed by a substance as a function of the wavelength. Different types of absorption spectroscopy result from different types of radiation used (i.e., X-ray, UV, visible, infrared). The absorption of visible light by a substance results from the excitation of electrons to higher energy levels. Because electron energy levels are quantized, different substances absorb different colors of visible light.

The fine line absorption spectra of the rare earth elements are indeed “rare.” They result from electron transmitions involving the f orbitals.

References

Special thanks to John Brodemus, Oswego, Illinois, for providing us the idea for this activity.

Hewitt, Paul G. Conceptual Physics, 3rd Ed.; Addison Wesley Longman: Menlo Park, California, 1999; pp 422–431.

Sadler, Philip. “Projecting spectra for Classroom Investigations”; The Physics Teacher; October 1991; pp 423–427.