Experimental Spectroscopy

Today, you have a chance to see different kinds of spectra for yourself, with your very own eyes. Simply mark your answers on these sheets in the spaces provided, and turn in the sheets at the end of class.

1. First, use the spectroscopes to observe a number of different types of light. Be careful. Make sure that you record the locations of at least the 4 strongest lines or features in each spectrum.

Observe, and sketch on your spectrum sheets, the spectrum of each of the following:

An Incandescent (normal light bulb) light: this is called a continuous spectrum. Write down the range of numbers that corresponds to each color.)
A Fluorescent light (carefully note the positions of the bright lines that appear in addition to the continuous spectrum.)
A Neon tube
A Hydrogen tube
A Mercury tube
A Helium tube
A Low-pressure Sodium lamp
A High-pressure Sodium lamp. Most of the RIT walk lights and parking lot lights are this sort of lamp. Look very carefully at the yellow-orange region of the spectrum.
In addition, observe and sketch the spectra of three other lights such as exit signs, TV or computer monitor screens (state the color of the region you observe; suggestion: look at a white section of the screen), colored bulbs (e.g., traffic lights, "neon" signs), etc.
Attach your spectra sheets to this handout.

2. Record the colors of the individual bright lines you observed in the fluorescent light and the readings on the scales that correspond to those lines below.
```
                reading in eV          reading in nm
  color          (top scale)          (bottom scale)
-----------------------------------------------------------------
```
Which scale (nm or eV) do you prefer to use? Why?
How do the lines you see in the fluorescent light compare to the lines you saw for the mercury tube? They should be very similar; this is how we know that there is mercury in standard fluorescent lights.
Compare the high-pressure sodium light spectrum with the low-pressure light spectrum. What do you see? In particular, discuss the yellow-orange part of the spectrum near 590 nm.
3. (a) Point the slit of your spectrometer at a white surface (like the screen) that has fluorescent light shining on it. Describe the spectrum you see and note the positions of any bright lines.
(b) How does the spectrum of the light reflected from the screen (or other white surface) compare to the spectrum of the overhead fluorescent lights? These should be very similar because the light from the fluorescent bulb is reflecting off the white surface.
4. (a) Look at a bright light bulb (or out the window at a bright sky, not at the sun). Now hold a bottle of transparent colored liquid in front of the spectrometer slit and notice how the spectrum changes. Record below, for at least 4 different bottles), the label on the bottle, the missing color(s), and the position(s) of the dark band(s) on the spectrometer scale. Be certain that at least one of your bottles has more than one missing color. Circle the units you use (eV or nm, depending on which scale you use).
```
   Label         missing color           position (eV or nm)
-----------------------------------------------------------------
```
(b) For the one of your bottles that has more than one missing color, compare the positions of the missing colors to the other positions that you measured for your other bottles. Do any of them agree? Try to find someone else in the class that has a matching missing color: give that person's name and the information about the bottle that she/he observed. This might mean that the bottle with two missing colors is a combination of two materials that each absorbs one color. What other bottles might have been mixed to make the one that you observed with more than one missing color?
Now visit two web sites and try to model what we have observed.
- Go to the MiniSpectroscopy site at Harvard:
  http://mo-www.harvard.edu/Java/MiniSpectroscopy.html
  1. Pick one element for which you observed the emission spectrum in part 1. For "your element", use the black-and-white graph in the middle of the applet to draw the spectrum you observed. Print your best result. Remember, brighter is taller on the graph.
  2. Choose the Sun as the "Source". Notice how subtle (easy to miss!) the "dark" lines are. Measure the locations of 5 of the absorption lines in the visible (colored, not grey) part of the spectrum. Do any of these match emission lines that you measured in part 1? What element(s) might be present in the sun based on your observations?
```
		   









	     
```
  3. Where do you think elements could be located in order for them to cause dark lines in the spectrum of the sun? List the one(s) of the choices below that are possible, and explain why you chose the answer(s) you did.
```
   inside the sun                    on the sun's surface
   above the sun's surface           in space between the sun and the earth
   in the earth's atmosphere 







        
```
- Go to the Virtual Quantum Mechanics site at Kansas State Univ.:
  http://phys.educ.ksu.edu/vqm/index.html
  Choose "Spectroscopy Lab Suite -- Gas Lamps -- Emission" to begin.
  1. Put the gas tube for "your element" into the power supply gas lamp sockets.
  2. Compare the spectrum that the computer shows to what you saw. Are there the same number of lines? Which shows more? Are the brightnesses of the lines similar to those you observed?
```
		
```
  3. Make a model of energy levels that matches your spectrum. (Follow the instructions at the bottom of the web page.) Note: to pick up and move a line, your mouse must be left of the vertical line; to draw a path showing an electron moving from one energy level to another, your mouse must be to the right of the vertical line; your mouse must be precisely placed in all cases.
  4. Print your result.
  5. If another lab group has the same element, compare your energy levels with theirs. If no other group has the same element, look at how groups approached their own element and create a second version of the energy levels for your element. Print it.
  6. Discuss how they are similar/different. Is one better or worse than the other? Why?
```
		 
```
  Now, switch to "Spectroscopy Lab Suite -- Gas Lamps -- Absorption".
  1. Put the gas tube for "your element" into the power supply gas lamp sockets.
  2. Compare the spectrum that the computer shows to what you saw and to what the computer showed for emission. Are there the same number of lines? Which shows more? Are the brightnesses of the lines what you observed?
```
		 
```
  3. Make a model of energy levels that matches your spectrum. (Follow the instructions at the bottom of the web page.)
  4. Print your result.
  5. If another lab group has the same element, compare your energy levels with theirs. If no other group has the same element, look at how groups approached their own element and create a second version of the energy levels for your element. Print it.
  6. Discuss how they are similar/different. Is one better or worse than the other? Why?
```
		 
```
Last modified 9/27/2001 by MWR
Copyright © Michael Richmond. This work is licensed under a Creative Commons License.