SCI : Light and Color - Case Notes

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Author - Dale Moore - Summer 1998

Background I've used an abbreviated version of this case as a one-day activity.  I suspect that the entire case as written would require at least two days, including one day with access to a flame source and both days with access to materials from the Chemistry Stockroom.

Teaching Notes

I. Goggle boxes

Materials: (Obtained from Chemistry Stockroom; please call Dale Moore at least one week in advance.)  The goggle boxes are small cardboard boxes, the kind that safety goggles come in, sealed with black tape.  A narrow slit is left open in one end, a small square window is in the opposite end.

Use: If the boxes are passed out to the class without any additional instructions, the students will discover through their own exploration that by looking at the room lights through the square window in the boxes (with the slit directed at the lights) that they can see a "rainbow" spectrum projected (apparently) inside the boxes. (The windows are diffraction gratings made from polarizing film. Diffraction can be loosely described as the "flaring out" of light as it passes through narrow apertures. In this case, the apertures are in the polarizing film. This flaring out occurs to a different extent according to the color of light being diffracted, thus resolving multicolored light into its components, the rainbow spectrum.) On further examination, students will begin to note that the rainbow spectrum that they're observing isn't a continuous spectrum, but rather has some more intense lines. At this point, they may be introduced to the elemental samples.

II. Elemental light bulbs

Materials: (Obtained from Chemistry Stockroom; please call Dale Moore at least one week in advance.)  The elemental samples are tubes containing vapor; we have samples of hydrogen, helium, neon, and mercury.  When the tubes are put into the power source, they're like a light bulb in a socket.  Still using goggle boxes.  {JS: a few example spectra and an explanatory diagram are also available here.}

Use: The students will now automatically turn their goggle boxes on the newly introduced light source.  Turning out the lights will increase visibility.  Going from hydrogen (the simplest of atoms with three visible color lines in its spectrum) to helium to neon and finally to mercury, students will notice an increase in the number of lines in the spectrum, and at the very least should notice that the pattern of lines that they observe is characteristic of the sample under examination.

Discussion: At this point the students have learned enough about spectroscopy to discuss its utility. In chemistry, we use atomic emission spectroscopy (AES) to identify and indeed to quantify elemental composition of samples. This technique works just like the students' goggle box experiment, a sample is energized (typically with a flame) and the resulting light is examined. This can be used, for example, to detect harmful materials in our environment. If a scientist observed the same pattern of light from an energized water sample that the students observed in looking at the mercury light bulb, then she would conclude that the water contained mercury, a dangerous contaminant.

Also, astronomers typically observe the same spectrum when looking at stars as the students did when looking at the hydrogen light bulb. Obviously, because stars have a large hydrogen content. In addition, scientists like Edwin Hubble (see Trefil & Hazen, CH15) noticed that when they looked at far-away galaxies, the hydrogen spectrum that they observed had a similar pattern of lines, but with the lines shifted (not at exactly the same colors). Hubble theorized that this was due to the movement of the galaxies with respect to our own galaxy caused by the expansion of our universe, quantifying this theory with an equation.

III. Flame tests

Materials: (Obtained from Chemistry Stockroom; please call Dale Moore at least one week in advance.)  Salt samples of elements with distinct color signatures, esp. Na (yellow), Cu (blue), Sr (red), and Ni (green); wires for dipping; and a burner with matches. Because the burner requires a gas source, this experiment must be done in Willet, or else an alternate flame source must be found. Still using goggle boxes.

Use:  The flame tests are essentially simplified versions of atomic emission spectroscopy (AES). By dipping the wire in a salt or salt solution sample, then holding the wire in the flame, a specific color is generated. The color observed is a function of the element being energized. (A particularly curious student may ask about the source of the color. Without much chemistry background, it's difficult to explain in any detail, but if you've talked about the atomic model, it may be easier. When energy goes into an atom, its orbiting electrons move to higher energy orbits - like going to higher altitude. Then, when the electrons fall back down again, the energy leaves the atom as light. The color of light is a function of the particular atom. It can be noted that this behavior was very important in the creation of an atomic model, as the model had to be consistent with the observation, i.e. it had to provide an explanatory mechanism by which the atom could absorb and then re-emit energy.) As with AES, the function of the flame tests is elemental identification, though a flame test is difficult to quantify.

IV. Nickel ion solutions

Materials: (Obtained from Chemistry Stockroom; please call Dale Moore at least one week in advance.) Solutions of nickel sulfate (green), ammonia, and ethylene diamine (all aqueous solutions, 1.0M in concentration). Caution: The ammonia and ethylene diamine solutions are basic and can cause irritation; avoid contact and wash after use. Also need clean beakers for mixing. Alternate discussion requires goggle boxes and light source (small desk lamp).

Use: Adding the ammonia to the nickel will turn the green solution to blue. Adding the ethylene diamine solution to either the blue or green solution will turn it violet.

Discussion: The colors that students observe in these solutions have different origins than in the previous demonstrations. In the case of the excited elements, they were emitting light. Observed colors were emitted colors. In the case of this material (and make the students figure this out), the samples are absorbing some colors, but not others. The observed color is that which hasn't been absorbed. In my class, I've emphasized this point with a color wheel and discussion of complementary colors. Materials tend to appear to be the color that complements the color being absorbed. (Red-absorbing materials appear green, orange-absorbing materials appear blue, yellow-absorbing materials appear violet, etc.)

We also discuss the energy of light (from low energy to high: red-orange-yellow-green-blue-violet), and I ask students to look for the trend (pattern) demonstrated by the nickel compound series (green to blue to violet). Ultimately, the trend is for materials combined more strongly (held together by stronger chemical bonds) to absorb light with higher energy. (Again, if you've talked about the atomic model, you can connect this to the fact that with stronger chemical combinations, the electrons are held more tightly; therefore, it takes more energy to move them to higher altitude (higher energy orbits).

Alternate Discussion: [New idea.] The difference between emission and absorption as sources of observed colors can also be demonstrated by using the colored solutions as light filters. Placing the colored solutions (in flasks or beakers with sufficient volume to block a light source) between a light source and a viewer armed with a goggle box allows that viewer to see a decrease in the intensity of particular colors coming from the source. This experiment would allow the students to construct the theory of complementary colors without any prior knowledge. Using a small light continuous source, it's somewhat difficult to single out particular colors; instructors may have to provide cues. Discussion could then be continued as above.