Introduction Experiment #9. Atomic Emission Spectroscopy Spectroscopy is the study of the interaction of light with matter. This interaction can be in the form of the absorption or the emission of electromagnetic radiation. When elements or compounds are exposed to large amounts of energy in the form of heat, light or electricity, they may absorb this energy. When energy is absorbed electrons can jump from their ground state, or lowest energy level, to an excited state, or higher energy level. Electrons in excited states are unstable and will eventually release energy again to return to lower energy states. This release of energy is what we observe in atomic emission spectra. A basic principle of quantum theory states that electrons can only have certain specific energy levels. Hence, when electrons move from one energy level to another, a specific amount of energy (a quantum) is released or absorbed. The amount of energy in any form of radiation is directly proportional to its frequency (E = hν), so the energy emitted when an electron moves to a lower energy state will have a distinct frequency and wavelength. Taken together, all of the wavelengths of light emitted from a particular atom from these electron movements constitute that atom's emission spectrum. Each element or compound has a distinct emission spectrum that can be used to help identify it. Atomic emission spectra can be thought of as atomic fingerprints. When a high electrical potential is applied to a tube of hydrogen gas, the atoms will absorb some of the energy and reemit it as light. The distinct wavelengths emitted appear as lines when viewed through a spectroscope. Hydrogen emits light in the infrared, visible and ultraviolet regions. The lines in the visible region, which correspond to electrons dropping from higher energy levels to n=2, are known as the Balmer series. You will observe the lines of the Balmer series in this lab. Hydrogen also emits wavelengths in the UV region, known as the Lyman series, when electrons drop to n=1, and in the infrared, known as the Paschen series, when electrons drop to n=3. The energy levels of the hydrogen atom are schematically represented in the diagram below.
Calculations Involving Energy Levels of Hydrogen The energy of the electron in a hydrogen atom is given by equation (1) where n is the quantum number of the energy level (n= 1, 2, 3,...) The energy difference between two energy levels in hydrogen is given by equation (2) where n f is the final energy level of the electron and n i is the initial energy level of the electron. This can be simplified to equation (3). (1) (2) (3) When n f > n i, ΔE will be positive energy is absorbed when an electron goes to a higher energy level. When n f < n i, ΔE will be negative energy is released when an electron drops to a lower energy level. The energy difference between levels (ΔE) is equal to the energy of the photon absorbed or emitted (E photon ). The energy of a photon is calculated with equation (4) where Planck s constant (h) = 6.626 x 10-34 Js, and the speed of light (c) = 3.00 x 10 8 m/s. Once we measure the wavelength of light in the hydrogen spectrum, we can use equation (4) to determine the energy of the photons. Since we are observing hydrogen s emission spectrum, we must use the negative E photon value for ΔE in equation (3). (5) We know that for the Balmer Series (the visible wavelengths of emitted light that will be observed in today s lab) n f = 2. We can use equation (3) to calculate the initial energy level (n i ) that the electron dropped from. In part B of the experiment, you will measure the wavelengths emitted by hydrogen atoms and ultimately determine which energy level transition it corresponds to. Every element has a distinct spectrum which can be used to identify it, much like a fingerprint. Helium was discovered when scientists looking at light from the sun noticed an absorption spectrum pattern that didn t correspond to any known element. In part A of the this experiment, three lamps containing unknown" gaseous atoms will be analyzed. You will determine their identity by comparing your observed wavelength values to a table of known values. Note that the Bohr equation (3) only applies to hydrogen, so we will not calculate energy levels of other elements. (4) Useful Physical Constants and Conversion Factors Planck's constant h = 6.626 10 34 Js 1 angstrom = 10 10 meter speed of light c = 3.00 10 8 m/s 1 nm = 10 9 meter
Laboratory Activity Equipment handheld spectroscopes, spectroscopes, 5000 volt transformer, lamps containing H, He, Hg, and Ne. Safety Hazards - DO NOT TOUCH THE LAMP OR THE METAL CONNECTIONS WHILE THE APPARATUS IS ON! Instructors: Please set up each apparatus for all the lamps. Place each power supply on one of the lab jacks and adjust the height of the lab jack so that the slit on the spectroscope lines up with the center of the lamp. Insert the lamp carefully in the power supply being careful not to touch the middle of the lamp. Only handle the ends of the lamp. Students: Part A 1. Hold a plastic handheld spectroscope to the fluorescent lights in the room, then to sunlight coming through the window. Align the slit with the brightest part of the light for the best results. Describe what you observe. 2. Turn on one of the unknown lamps for up to one minute. You may touch the sides of lamp to steady it as you flip the switch, but do not touch the center of the lamp or the electrical connections. After one minute, turn the lab off and allow it to rest for at least one minute before turning it on again if necessary. 3. Look through the spectroscopes in front of the lamps for unknowns #1, #2, and #3. Record the wavelengths of the major lines to two significant figures. Note the unit on the spectroscope. For some of the unknowns, you will see an almost continuous spectrum rather than discrete lines. Record the range and center wavelength of this broad band of color. Compare the wavelengths to those in Table 1 and determine the identity of each unknown. Part B 3. Use the same technique as in Part A to record the color and wavelengths of the hydrogen spectrum. You should be able to see three or four lines. Use these wavelengths to calculate the ΔE for the transition and then the initial energy level of the electron for each transition (n i ) Remember that n f = 2 for the visible region (the Balmer series).
CHM111 Lab Atomic Emission Spectroscopy Grading Rubric Name Team Name Criteria Points possible Points earned Lab Performance Printed lab handout and rubric was brought to lab 3 Safety and proper waste disposal procedures observed 2 Followed procedure correctly without depending too much on instructor or lab partner 3 Work space and glassware was cleaned up 1 Lab Report Observations and data recorded with proper units 1 Correct identification of unknowns 2 Calculations shown clearly and completely with units. 4 Question 1 1 Question 2 (calculations shown in detail with units) 2 Question 3 (calculations shown in detail with units) 1 Total 20 Subject to additional penalties as per the instructor
Atomic Emission Spectroscopy: Data Sheet Name Part A: 1. Record observations: fluorescent light sunlight Based on what you observed, describe is the difference between a continuous spectrum and a line spectrum. 2. Record color and wavelength of spectral lines for the unknowns. Include the correct units with the wavelength. Unknown # 1 Unknown #2 Unknown #3 Color Wavelength Color Wavelength Color Wavelength Use the following table to dentify each of the unknown elements: CHARACTERISTIC SPECTRAL LINES * He (nm) Hg (nm) Ne (nm) 447 405 540 502 436 583 588 546 585 668 579 640 707 Unknown #1 Unknown #2 Unknown #3 * This is an abbreviated table, but it shows sufficient wavelengths for identifying the unknown elements. Longer tables are in the HANDBOOK OF CHEMISTRY and PHYSICS Report Page 1 of 3
Atomic Emission Spectroscopy: Post Lab Name Part B: 1. Record color and wavelengths for the hydrogen spectrum, then use equations (4), (5) and (3) to calculate ΔE and n i. Color wavelength E photon n f n i (make ΔE negative, use separate page if needed) 2 2 2 2 Report Page 3 of 3
Atomic Emission Spectroscopy: Post Lab Name 1. Rank the following radiations from lowest to highest energy. Radio waves Infrared waves Gamma rays microwaves X rays Lowest E Highest E 2. Which of the following electrons will emit light of LONGER WAVELENGTH? An electron dropping from n = 3 to n = 2 OR an electron dropping from n = 4 to n = 3? Calculate the wavelength for each transition to justify your answer. 3. The ionization energy is the energy needed to remove an electron from an atom which corresponds to a raising the electron from n=1 to an orbit that has n=. What is the energy needed to remove the electron from a hydrogen atom? Report Page 3 of 3