Virtual Atomic Spectra NAME: INTRODUCTION Atoms have multiple discrete electron energy levels. When electrons in gas phase atoms are energized, either by an electric field or by heating, they can be excited to a higher energy level within the atom. When the source of the excitation is no longer present, the electrons “relax” back R G B V down to lower energy levels by releasing energy in the form of a photon of light. Consequently, photons of different colors of light can be emitted from the same element, producing an atomic spectrum. The figure to the right shows the atomic spectrum produced by hydrogen atoms. The electron can be excited from the ground state (n=1) to any of these excited states. The energy difference between the energy levels produces a unique color of light. For example, the 656 nm light (red) is emitted from a hydrogen atom when an electron relaxes from the second excited state (n=3) to the first excited state (n=2). PROCEDURE 1. Open Virtual Chem Lab and go to the Quantum module. Once in the lab, go to the stockroom. Double click on the gas holder first, then get some H 2 from the gas canisters. You can’t pick the gas until you take out the gas holder. Double click on the spectrometer (in the detectors section) and the electric field (in the modifiers section). 2. Go back to the lab and set up by placing the spectrometer in the right center location of the lab bench. Place the gas sample in the middle of the table and the electric field directly on the sample. Turn the electric field up to 300 by clicking on the tiny arrows on the upper part of the electric field. Turn on the spectrometer by clicking on the green button. Set the switches to “frequency” and “visible.” 3. Record your observations. Repeat these steps for Neon (Ne) and Mercury (Hg). If you leave the spectrometer and the electric field behind, you can switch out gases without having to reset the whole lab. 4. Close the Quantum module and open the Inorganic module. Double click on a test tube to get one. Click on one of the metal ions given in the table below (sodium, barium, strontium, and copper). Then, click on the Bunsen burner at the bottom left to do a flame test. Choose the Bunsen burner without the blue cobalt screen. Click the red container to the right to clear the lab between metal ions. 5. Watch what happens. Wait a few seconds for the color of the flame to change. Record what the final color of the flame is. Repeat for the other ions. Gas in the tube Colors of lines in spectrum H2 Color seen with eye Number of lines observed using spectrophotometer (doesn’t need to be exact, can be too many to count) Purple 3 Ne Red Too many to count All colors Hg Blue Like 50 plus All colors Substance in flame Strontium Sodium Flame color Pink purple Orange yellow Substance in flame Barium Copper Red, green, and blue Flame color Yellow reddish Blue greenish QUESTIONS 1. A fluorescent light bulb contains hot gaseous mercury atoms. An incandescent bulb contains a thin piece of metal called a filament. It is heated until it glows, a phenomenon known as blackbody radiation. Based on what you know about atomic spectra, will the fluorescent bulb or the incandescent bulb have a continuous spectrum? Explain. Filament = continuous spectrum because of blackbody radiation. Fluorescent = discontinuous spectrum because it has hot atoms that are creating different emission and absorption spectra. If you need help deciding, you can go into the virtual module. Go into the stockroom, click on the clipboard, and select the “Blackbody radiation” experiment. Flip the switch to “visible” and compare what you see to what you observed earlier with the hot gaseous elements. 2. Imagine that you are designing custom fireworks. Which metal ions that you did the flame tests on (not the gases you used to look at spectra) would you use to create the following colors of fireworks? You may use one metal ion or more than one metal ion for each box. Although other metal ions are possible, you can create all of these colors with what was used in the flame tests above. Firework Color Metal Ion(s) Orange Sodium Red Copper and Strontium Purple Strontium 3. Each element has its own distinct emission color bands. Why are the emission spectra of no two elements the same? They aren’t the same because there are a different number of electrons in each element which results in differing color bands. 4. The speed of light is 2.99792 x 10 8 m/s. All wavelengths of light travel at the same speed. The wavelength associated with one of the colors of light emitted by hydrogen is 486.133 nm. Convert this wavelength to frequency using this relationship: c= νλ, where λ is wavelength, ν is frequency, and c is the speed of light. Speed of light = frequency x wavelength Frequency = speed of light / wavelength ((2.99792 x 10-34 m/s) x 109 nm/m) / 486.133 nm = 6.166 x 1014 s-1 or Hz 5. The energy associated with a photon of light energy (in units of Joules, J) can be calculated using the relationship E=hν where E is energy, h is Planck’s constant (6.62607 x 10 -34 J·s) and ν is the frequency of the light energy. This is also the energy released as an electron relaxes from a higher energy state to a lower energy state. Calculate the energy associated with a photon that has a wavelength of 486.133 nm. E = hv (6.026 x 10-34) x (6.166 x 1014) = 4.086 x 10-19