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CHEMISTRY SEMESTER ONESPECTROSCOPY LAB

EMISSION SPECTROSCOPY

Lab format: this lab is a remote lab activity

Relationship to theory: This activity covers the relationship between colors and absorbed/emitted light, as well as the relationship between absorption of electromagnetic radiation and global warming.

Instructions for Instructors: This protocol is written under an open source CC BY license. You may use the procedure as is or modify as necessary for your class. Be sure to let your students know if they should complete optional exercises in this lab procedure as lab technicians will not know if you want your students to complete optional exercise.

Remote Resources: Primary - UV/Vis Spectrometer; Secondary - Emission Tubes

Instructions for Students: Read the complete laboratory procedure before coming to lab. Under the experimental sections, complete all pre-lab materials before logging on to the remote lab, complete data collection sections during your on-line period, and answer questions in analysis sections after your on-line period. Your instructor will let you know if you are required to complete any optional exercises in this lab.

ContentsEMISSION SPECTROSCOPY.............................................................................................1

Learning Objectives......................................................................................................2Background Information...............................................................................................2Equipment....................................................................................................................3Preparing to Use the Remote Web-based Science Lab (RWSL).................................3Introduction to the Remote Equipment and Control Panel...........................................3Experimental Procedure...............................................................................................4Appendix: Introduction to the Remote Equipment and Interface.................................9

Creative Commons Attribution 3.0 United States License 1


LEARNING OBJECTIVES

Identify the most intense peaks in the emission spectra for several molecular and atomic gases.

Use the visible emission spectrum for hydrogen gas to determine the electronic transitions taking place and compare these with the theoretical predictions using the Rydberg equation.

Relate wavelength and color of light. Be able to calculate frequency and energy of electromagnetic radiation from its known

wavelength. Be able to predict absorption of light by a gas from its emission spectrum Identify an unknown gas by its emission spectrum

BACKGROUND INFORMATION

Spectroscopy is the study of the interaction between light and matter, in particular how atoms or molecules absorb or emit electromagnetic (EM) radiation. Electromagnetic radiation is characterized by its wavelength or its frequency, which are related by the equation c = λν, where λ is wavelength (typically units are nanometers for visible light), ν is frequency (units are s -1, known as Hertz or Hz) and c is the speed of light (299,792,458 m/s). Note that for any wave, its wavelength times its frequency equals its speed; for electromagnetic radiation that speed is the speed of light. Also note that as wavelength increases, frequency decreases, and vice versa.

Electromagnetic radiation may be understood as having both wave-like and particle-like properties, depending on the experimental setup and the type of detector used (in this experiment, we are measuring wavelengths using a spectrometer that emphasizes the wave-like properties of light). From a particle-like point-of-view, the smallest amount of EM radiation (of a particular wavelength) that can be emitted or observed is called a photon. As Albert Einstein showed more than a century ago, the energy of a photon is proportional to the frequency of the EM radiation involved (and thus inversely proportional to the wavelength): Ephoton = hν, where h is called Planck’s constant (h = 6.62607 x 10 -34 J·s). Since energy is conserved, and since atoms and molecules almost always emit or absorb one photon at a time, the change in energy of an atom or molecule will determine the frequency (and wavelength) of the EM radiation absorbed or emitted: ΔEatom = Ephoton = hν = hc / λ.

Atoms (or molecules) emit photons of very distinctive wavelength (as determined by the previous equation) when the energy of the atom decreases. Atoms can absorb energy when they absorb a photon with an appropriate wavelength (a photon whose energy matches the difference in energy between the initial and final energy states of the atom, ΔEatom). Since atoms have quantized energy levels (only certain energy levels are possible), there are only certain values possible for ΔE atom. Thus, only certain frequencies and wavelengths of EM radiation will be emitted by the atoms (and molecules) in this experiment.

For most atoms, determining the energy levels possible involves very complex quantum mechanical calculations. The H atom (hydrogen) is a notable exception, since it is the simplest of atoms, with only a single electron. Even before quantum mechanics was discovered and understood, physicists



and chemists were able to write a very simple equation to predict the wavelengths for the EM radiation absorbed or emitted by H atoms, known as the Rydberg equation: 1 / λ = RH {(1/n12) – (1/ n22)} where RH is called the Rydberg constant (0.010967760 nm-1), and n1 and n2 are positive (non-zero) integers called quantum numbers, such that n1 < n2. See your text for more details about the Rydberg equation and to see an example of using this equation in a calculation.

You may notice that we are using hydrogen molecules (H2) in this experiment, whereas the Rydberg equation only applies to hydrogen atoms (H). Hydrogen atoms are not chemically stable, but H 2

molecules are stable and can be used to fill the glass tube in the hydrogen lamp. The high-voltage electrical discharge used to power the lamp will temporarily break up H2 molecule into individual H atoms, which then gain additional energy from the discharge; this additional energy is then emitted in the form of visible light as the H atoms return to lower energy states.

The energy of a photon of electromagnetic radiation is given by the relationship: E = hν where E = energy in joules, ν = frequency in cycles per second, and h = Planck’s constant = 6.62607 x 10 -34 J·s

The relationship between wavelength and frequency of electromagnetic radiation is: λν= c where λ = wavelength in meter and c = 2.996 x 108 m/s, the speed of radiant energy in a vacuum

EQUIPMENT

Paper

Pencil/pen

Computer with Internet access

PREPARING TO USE THE REMOTE WEB-BASED SCIENCE LAB (RWSL)

Click on this link to access the InstallGuide for the RWSL: http://denverlabinfo.nanslo.org Follow all the directions on this webpage to get your computer ready for connecting to the remote lab.

INTRODUCTION TO THE REMOTE EQUIPMENT AND CONTROL PANEL

Watch this short tutorial video to see how to use the RWSL control panel: http://denverlabinfo.nanslo.org/video/emission_spectroscopy.html

There are appendices at the end of this document that you can refer to during your lab if you need to remind yourself how to accomplish some of the tasks using the RWSL control panel.


http://denverlabinfo.nanslo.org/video/emission_spectroscopy.html

http://denverlabinfo.nanslo.org/


EXPERIMENTAL PROCEDURE

Read and understand these instructions BEFORE starting the actual lab procedure and collecting data. Feel free to “play around” a little bit and explore the capabilities of the equipment before you start the actual procedure.

Once you have logged on to the Remote Lab, you will perform the following Laboratory procedures:

1. Use the control panel to gather data from the emission lamps.a. Be sure to start the spectrometer so you can view the spectra when the lamps

are energized.b. Use the camera to zoom in on each emission lamp to read the labels and

determine what gas is in each one.c. Use the screw-drive robot to position the fiber optic cable for the spectrometer

and record the spectrum of each of the emission lamps. d. While viewing each spectrum, use the cursor to find the wavelength of five or six

most intense peaks for each of the gases. Record these in a table.e. Also, while each lamp is glowing, zoom in close with the camera and see what

color it appears to be.

2. Use the Rydberg equation (refer to your text) to predict the wavelength of the electromagnetic radiation emitted for following electron transitions for the hydrogen atom (rounded-off to the nearest 0.1 nm).

Electron Transition Predicted Wavelength (nm)

n = 3 → 2

n = 4 → 2

n = 5 → 2

n = 4 → 3

n = 2 → 1

3. Using Figure 1 and your predicted wavelengths above, determine what type of electromagnetic radiation is produced by these electron transitions in the hydrogen atom. In other words, does the transition of n = 3 → 2 produce ultraviolet (UV), visible or infrared (IR) radiation? What about the other transitions above?



Figure 1 - By Jonathan S Urie [CC-BY-SA-3.0 (http://creativecommons.org/licenses/by-sa/3.0)], via Wikimedia Commons

Electron Transition Type of Radiation

n = 3 → 2

n = 4 → 2

n = 5 → 2

n = 4 → 3

n = 2 → 1

4. How closely do the first three transitions listed in table above correspond to the observed wavelengths of the three largest peaks in the hydrogen atom emission spectrum?

5. Note that all the most intense peaks for neon gas have wavelengths greater than 580 nm. Based on Figure 2, what colors of visible light are emitted by the neon tube? Does this explain the apparent color of neon lamps to the ‘naked’ eye?



Figure 2 - Visible Portion of EM Spectrum

6. Shown below is a portion of the emission spectrum produced by a mixture of two of the gases involved in this experiment. Based on your experimental results, does this gas mixture include helium gas? Explain your reasoning. Can you determine which two gases are in this gas mixture?

350 400 450 500 550 600 650 7000

2000

4000

6000

8000

10000

12000

14000

Wavelength (nm)

Inte

nsith

y (c

ount

s)

Figure 3 - Spectrum of Gas Mixture

7. Iron vapor produces an emission spectrum that includes an intense peak at 527.0 nm. Determine the frequency (in Hz) for this type of electromagnetic radiation. What color of visible light corresponds to this wavelength? What is the energy (in J) per photon emitted at this wavelength? What is this photon energy in units of kJ per mole? (In other words, one mole of these photons with wavelength of 527.0 nm consists of how many kilojoules of electromagnetic energy?)

8. The peaks that you observe in the emission spectrum of each gas are also wavelengths of light that the gas will absorb better than others. So, if a gas shows an emission peak at 550 nm, the gas will also absorb light with a wavelength of 550 nm. The more intense the emission peak, the more that light will be absorbed by the gas. Based on this


750 nm 380 nm450 nm500 nm580 nm600 nm620 nm


information, rank the four gases you observed in this experiment in order of how well they will absorb infrared light. Write the strongest infrared absorber on the left and the weakest on the right:

Best absorber > next best > next best > worst absorber

9. Based on these results, and the reference listed below, why do you think carbon dioxide is considered a “greenhouse gas” that we need to be concerned about, compared to the other gases you observed in this experiment?

Reference: http://www.eoearth.org/article/Atmospheric_composition (Bear in mind that these results only take a small portion of the infrared portion of the spectrum into account.)

10. Figure 4, below, shows the absorption spectra of several common gases that are prevalent in the atmosphere. Peaks in this spectrum indicate the wavelengths that these gases absorb the best. The horizontal axis is the wavelength in microns, which is another name for micrometers. Note that this horizontal axis is a logarithmic axis, so this single figure covers a large portion of the electromagnetic spectrum, namely the ultraviolet, visible and infrared regions of the spectrum. The vertical axis is the percent absorptivity of each gas (in other words, what percent of EM radiation at a given wavelength is absorbed by each gas). A percent absorptivity of zero means a gas is completely transparent at that wavelength; a value of 100 means complete absorbance at that wavelength, so the gas is opaque. The spectrum labeled Total corresponds to the spectrum of the total atmosphere.

a. The visible spectrum is typically defined as wavelengths between 400 and 700 nm. Convert these wavelength values to microns and locate the visible region in the figure below. Also locate the ultraviolet and infrared regions of the figure (hint: does infrared radiation have longer or shorter wavelengths than visible light?).

b. Does the spectrum labeled Total indicate whether the atmosphere is mostly transparent or mostly opaque in the visible region? Briefly explain your answer.

c. The figure also shows that the Earth emits large amounts of “thermal radiation” at wavelengths between about 5 and 50 microns; is this in the ultraviolet, visible or infrared region of the spectrum? Greenhouse gases can cause global warming by absorbing this emitted “thermal radiation”, trapping heat in the Earth’s atmosphere. Are carbon dioxide, water or oxygen (plus a little ozone, O3) greenhouse gases according to this figure? Briefly explain your answer.

d. Why are scientists and governments much more concerned about carbon dioxide acting as a greenhouse gas than water or oxygen? What effect do human activities have on the levels of these gases in the atmosphere?


http://www.eoearth.org/article/Atmospheric_composition


Figure 4 - Solar Radiation

11. Identify the “unknown” gas in the emission lamp labeled “E”. Explain why you think you are correct.



APPENDIX: INTRODUCTION TO THE REMOTE EQUIPMENT AND INTERFACE

Activating the Spectrometer

1. When you first start the Emission Spectrometer control panel, you will see the Control Panel seen in Figure 5. You can gain control of the interface by right-clicking anywhere on the screen and selecting “Request Control of VI”. If someone else has control, you will get a message telling you that you need to wait. You will be placed in queue and will get control of the interface when the other person releases it.

Figure 5 - Emission Spectrometer Control Panel

2. Click the Start button on the far left portion of the control panel to activate the spectrometer. The button will now turn yellow and say “Pause”.

3. In the video window on the right side of the screen, you should see four black boxes labeled A through E, and also labeled (from left-to-right) H2, He, Ne, CO2 and ?. These are ‘emission lamps’ that generate different colors of visible light, depending on the gas that is in them. In front of these lamps is a horizontal screw-drive robot with a mounted



fiber optic cable that transmits the detected light to the spectrometer, where the spectrum will be digitally acquired. The spectrometer is the small box located near one end of the track.

4. The Camera Preset Positions allow you to zoom in quickly to one of the 5 lamps. Hovering over the gray area where the buttons are will give you a pop-up menu that describes where each Preset is assigned to. You can also zoom in even closer using the ZoomIn and ZoomOut buttons.

Figure 6 – Presets will zoom to various lamps

5. To acquire the spectrum of one of the emission lamps, first select Preset 6 so that you can see the sensor move along the track. Click on the green-colored A button on the left side of the screen. This will position the sensor in front of this tube and turn on the H2 lamp. NOTE: because the high voltage involved in these lamps, DO NOT set the TimeTubeON field to more than 120 seconds. The lamp status and the number of seconds it has been energized are shown in the Spectrum Status portion of the interface screen (to the right of the green lamp buttons A – E). Click on Preset 1 to zoom in on the H2 lamp once it is turned on.



6. Use the “Nudge Left” and “Nudge Right” buttons to move the fiber optic and make sure you are getting the maximum signal from the lamp. If the lamp turns off before you’re finished, just click the same letter again and it will re-energize. You can also change the time that the lamp stays on (in seconds) by changing the “TimeTubeON” field.

7. Once you have the signal maximized, click the “Pause” button to hold the image so the light doesn’t have to be energized while you locate the maximum peaks (Figure 7).

Figure 7 – Press the Pause button to “freeze” the spectrum

8. You will now need to “zoom out” on the spectrum window to view the entire spectrum properly. Here’s how to zoom in and out on the spectrum:

a. Click on the center button at the lower right of the graph, shown below in Figure 8.

b. This brings up a small sub-menu of other buttons. The only two that are useful to you are the left-most buttons in the top and bottom rows, although you can play around with the others if you want to. Select the left-most button in the bottom row to view the entire spectrum.



Figure 8 - Spectrum Zoom Out button

c. Select the left-most in the top row to select specific parts of the spectrum to “zoom in” on and view more closely (see Figure 9). After clicking this button, you use the mouse to draw a box around the area that you want to zoom in to. Be sure you draw the box so that it includes some area past the top of the peak you are interested in, or else it will chop off the top of it in the viewing window.

d. If you accidentally zoom in too far or on the wrong part of the spectrum, just zoom out and start over again.


Spectrum Zoom Out


Figure 9 – Spectrum Zoom In button

9. Export a graph of the spectrum: a. Locate the Export to Clipboard button, but don’t click on it yet. Go to the pull-

down box to the right of it and set it to Graph Image. Now click the Export to Clipboard button, which will place a copy of the spectrum in your clipboard.

b. Start a program like Paint or Word or Powerpoint and paste in the spectrum. Save the file with an appropriate name so you can find it later.

c. You must have control of the Control Panel to do this, of course.10. Using the cursor to locate the most intense peaks:

a. Make sure you are zoomed out to view the entire spectrum.b. Click the button labeled “Enable Cursor” under the left side of the graph – the

green light will come on and a vertical green cursor line will appear on the screen.


Spectrum Zoom In


c. Click the cursor control button indicated below in Figure 10.d. Use the mouse to “grab” the cursor line by clicking on it and dragging it to the

peak that you want to identify.e. If you want to zoom in on a peak for a closer look, make sure you place the

cursor approximately on that peak before you click the Spectrum Zoom In button and draw a box around it.

f. If you lose the cursor while zooming in on peaks, just zoom out again to find it.

Figure 10 – Enable Cursor and Cursor Control button

g. There are now two fields under the graph: “Wavelength (nm)” and “Intensity”. The Wavelength field shows the current position of the cursor, and Intensity shows a relative intensity reading of wherever the cursor is located (Figure 10).


Cursor Control


h. Use the cursor to find the wavelength of each major peak in the spectrum.i. Once you have the cursor on top of a peak, you can zoom in on it to make sure

you are really on the highest part of the peak. (Sometimes they are double peaks!) If you zoom in or out, you will need to click the cursor control button again in order to move the cursor.


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