CHEMISTRY SEMESTER ONE

Transcription

1 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. LEARNING OBJECTIVES 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 (begins on page 10). 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 BACKGROUND INFORMATION Spectroscopy studies the interaction of 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 = λ ν, 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 = h ν, where h is called Planck s constant (h = x 10- Creative Commons Attribution 3.0 United States License 1

2 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 = h ν = h c / 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 if the atom, ΔEatom). Since atoms have quantized energy levels (only certain energy levels are possible), there are only certain values possible for ΔEatom. 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 ( 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). H atoms are not chemically stable, but H2 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. EQUIPMENT Paper Pencil/pen Computer (access to remote laboratory) Creative Commons Attribution 3.0 United States License 2

3 PREPARING TO USE THE RWSL SPECTROMETER CHEMISTRY SEMESTER ONE Setting up your computer for use with the RWSL: Ensure that your computer system is capable of interacting with the RWSL microscope. Currently RWSL works only on the Microsoft Windows operating system (XP or later) and a relatively up-to-date browser. To confirm that your system meets minimum requirements, visit this website: and follow the steps provided. Scheduling time at the RWSL Go to your online class website in D2L and open the RWSL Scheduler. Select the date and time you would like to attend lab. Try to choose a classroom that already has students scheduled in it so you have some lab partners to work with. Before you connect to the RWSL spectrometer: Open the Mumble software and connect to the Denver NANSLO server. This will establish contact with the Laboratory Technicians in the lab so they can assist you if you have any trouble. Connecting to the RWSL spectrometer When it is time to attend your scheduled lab, go to your online class website in D2L and open the RWSL Scheduler. There will be a link just above the calendar that allows you to access the lab. NOTE: This link will not be available until the exact time that your lab activity starts. Figure 1 - RWSL Scheduler Link to Lab Creative Commons Attribution 3.0 United States License 3

4 INTRODUCTION TO THE REMOTE EQUIPMENT AND INTERFACE: CHEMISTRY SEMESTER ONE DO NOT BEGIN WORKING ON THE LAB PROCEDURE UNTIL YOU HAVE READ ALL OF THIS INTRODUCTORY SECTION. THE PROCEDURE BEGINS ON PAGE In the video window on the right side of the screen, you should see four black boxes labeled A through D, and also labeled (from left-to-right) H 2, He, Ne and CO 2. These are four different lamps that generate different colors of visible light, depending on the gas used. In front of these lamp boxes is a horizontal camera track 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 (again, black-colored) located behind the right end of the track. 2. Select Preset 1 to zoom in on the H 2 lamp (Presets 2 through 4 zoom in on the other corresponding lamps). The black lamp boxes contain glass or transparent quartz tubes filled with small amounts of one of the various gases mentioned above. These small glass tubes are located within the black lamp boxes, behind the small vertical opening in the top half of the box. The glass tubes have imbedded electrodes, and when a high voltage is applied to these electrodes, the atoms and molecules that make up the gas become energetically-excited, causing them to emit visible light, as discussed in the Introduction above. Figure 2 - Preset 1 zooms in on the hydrogen lamp Creative Commons Attribution 3.0 United States License 4

5 3. Preset 5 zooms in on the spectrometer itself, if you want to see it, and preset 6 zooms back out to view the whole assembly. You can also pan, tilt and zoom the camera using the keypad controls on the screen if you want to. 4. Now let s acquire the atomic emission spectrum of hydrogen gas, H 2. 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 H 2 lamp. NOTE: because the high voltage involved in these lamps, they become very hot; to prevent overheating, each lamp will automatically turn off after 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 D). Click on Preset 1 to zoom in on the H 2 lamp once it is turned on. 5. Click the green Start button on the far left portion of the control panel to activate the spectrometer. The button will now turn yellow and say Pause. Note the lamp is now glowing brightly; what color does it look to your eye? Record color of each lamp for the Discussion questions later. Figure 3 - Press the Start button and it becomes the Pause button 6. You will now need to zoom out to view the entire H 2 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 4. Creative Commons Attribution 3.0 United States License 5

6 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 (see Figure 5), although you can play around with the others if you want to. Select the leftmost button in the bottom row to view the entire spectrum. c. Select the left-most in the top row to select specific parts of the spectrum to zoom in on and view more closely. 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. Figure 4 - Spectrum manipulation button Creative Commons Attribution 3.0 United States License 6

7 Zoom In Zoom Out Figure 5 - These two buttons are the most useful ones 7. If the timer runs out while you are working, just wait 10 seconds or so, and then click the letter button that corresponds to the lamp you were working with again. The lamp will light up and you will be right where you left off on the screen. As you get more practiced with the controls, you won t need as much time to accomplish your tasks. 8. Use the nudge left and nudge right buttons to move the fiber optic cable left or right until you get the tallest peaks possible. This will ensure the best possible signal. It may be that the tallest peaks are chopped off at the top of the graph, meaning that they are too intense for the spectrometer to measure. If so, use the nudge left and nudge right buttons to move the fiber optic cable left or right until the tallest peaks are no longer chopped off. You won t see the fiber optic move very much when you do this. These buttons only move it 0.1 mm in either direction. 9. Export a graph of the spectrum: a. Click the yellow Pause button to freeze the observed spectrum. b. Locate the round green 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 small green Export to Clipboard button, which will place a copy of the spectrum in the clipboard. c. Minimize the Internet Explorer window and return to your computer s desktop. 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. Creative Commons Attribution 3.0 United States License 7

8 10. Each student needs to download the graph. Each of your lab partners must take control of the interface while the graph is still frozen and download the graph as in Step Locate the most intense peaks using the cursor (no more than 6 peaks): a. Make sure you are zoomed out to view the entire spectrum (see step 9). b. Click the button labeled 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. c. Click the cursor control button indicated below in Figure 6. 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 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 6 - 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 7). Creative Commons Attribution 3.0 United States License 8

9 Figure 7 - Cursor, Wavelength and Intensity 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. If you zoom in or out, you will need to click the cursor control button again in order to move the cursor. j. Write down the locations of the peaks because you will use this information later. k. You can ignore the Show Absorbance Spectrum button, as it is not used in this experiment. l. Use the Export to Clipboard button to copy either the graph or the data table for the graph (you can select which one) to the clipboard. Then, you can open a document and use the Paste function to paste it from the clipboard. This way, you can include the graphs or data tables in your lab report. 12. Complete this process for the remaining gases. Each student should take a turn collecting the spectrum for one of the gases. Creative Commons Attribution 3.0 United States License 9

10 EXPERIMENTAL PROCEDURE: (REFERENCE THE ABOVE SECTIONS FOR DETAILS) 1. Log into Mumble and establish communication with the Lab Technician. 2. Using the link in the RWSL Scheduler on your course webpage, access the RWSL and take control of the interface. 3. Make a table of the wavelengths of the five or six most intense peaks in the acquired spectra of each of the four gases utilized, except hydrogen. Since hydrogen (H 2 ) has a much simpler spectrum than do the other gases, you only need to report the wavelengths of the three most intense peaks for H 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 = Using Figure 6.4 in your textbook (Brown, et al.) 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? 6. 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? 7. Note that all the most intense peaks for neon gas have wavelengths greater than 580 nm. Based on Figure 6.4 of the textbook, what colors of visible light are emitted by the neon tube? Does this explain the apparent color of neon lamps to the naked eye? 8. 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? Creative Commons Attribution 3.0 United States License 10

11 Intensithy (counts) CHEMISTRY SEMESTER ONE Wavelength (nm) 9. Iron vapor produces an emission spectrum that includes an intense peak at 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 nm consists of how many kilojoules of electromagnetic energy?) 10. 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 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 11. 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: (Bear in mind that these results only take a small portion of the infrared portion of the spectrum into account.) 12. The figure 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 Creative Commons Attribution 3.0 United States License 11

12 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, O 3 ) 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? Creative Commons Attribution 3.0 United States License 12