Lab: Excited Electrons

Transcription

1 Part A: EMISSION SPECTROSCOPY Lab: Excited Electrons According to the Bohr atomic model, electrons orbit the nucleus within specific energy levels. These levels are defined by unique amounts of energy. Electrons possessing the lowest energy are found in the levels closest to the nucleus. Electrons of higher energy are located in progressively more distant energy levels. If an electron absorbs sufficient energy to bridge the gap between energy levels, the electron may jump to a higher level. Since this change results in a vacant lower orbital, the configuration is unstable. The excited electron releases its newly acquired energy and falls back to its initial or ground state. Often, the excited electrons acquire sufficient to make several energy level transitions. When these electrons return to their ground state, several distinct energy emissions occur. The energy that electrons absorb is often of a thermal or electrical nature, and the energy that electrons emit when returning to ground state is electromagnetic radiation. In 1900, Max Planck studied visible emissions from hot glowing solids. He proposed that light was emitted in packet of energy called quanta, and that the energy of each packet was proportional to the frequency of the light wave. According to Einstein and Planck, the energy of the packet could be expressed as the product of the frequency ( ) of emitted light and Planck s constant (h). E = h If white light passes through a prism or diffraction grating, its component wavelengths are bent at different angles. This process produces a rainbow of distinct colors known as a continuous spectrum. If, however, the light emitted from hot gases or energized ions is viewed in a similar manner, isolated bands of color are observed. These bands form characteristic patterns, unique to each element. They are known as bright line or emissions spectra. By analyzing the emission spectrum of hydrogen gas, Bohr was able to calculate the energy content of the major electron levels. Although the electron structure as suggested by his planetary atomic model has been modified according to modern quantum theory, his description and analysis of spectral emission lines are still valid. In addition to the fundamental role of spectroscopy played in the development of today s atomic model, this technique can also be used in the identification of elements. Since the atoms of each element contain unique arrangements of electrons, emission lines can be used as spectral fingerprints. Even without a spectroscope, this type of identification is possible since the major spectral lines will alter the color of the flame. In the following experiment, you will use a spectroscope to examine several continuous and bright line spectra. OBJECTIVES: 1. To understand the relationship between atomic structure and emission spectroscopy. 2. To use a spectroscope and observe several sources of continuous and bright line spectra. MATERIALS: Student spectroscope Low wattage incandescent light source Fluorescent light source High-voltage power supply (to be used by teacher only) Various spectral tubes (to be used by teacher only)

2 PROCEDURE 1. Obtain spectroscope and examine the spectrum emitted by an incandescent light source. Record your observations. 2. Aim the spectroscope at a fluorescent light source. Compare this spectrum with the spectrum observed in Step 1. Record your observations. 3. Since the high voltages required for this part of the experiment may present an electrical hazard, you should not touch the spectrum tube or high-voltage assembly. Your teacher will load the spectrum tube and assemble the high-voltage power supply. 4. Once the power supply is turned on, aim the spectroscope at the glowing tube. Record the position of the bright emission lines. 5. Repeat Step 4 for each of the gas samples and record your observations.

3 DATA AND OBSERVATIONS (Cut and tape each one into your lab notebook) 1. Incandescent spectrum (Color or write in the color that you see) Description: 2. Fluorescent spectrum (Color or write in the color that you see) Description: (Draw the spectral lines that you see) 3. Gas Tubes: Gas Tube Red Orange Yellow Green Blue Violet

4 QUESTIONS 1. According to Bohr s atomic model, where may an atom s electrons be found? 4. Compare the spectra produced by incandescent and fluorescent sources. 5. How is a continuous spectrum different from an atomic emission spectrum? Explain. 6. Prior to its discovery on Earth, helium s existence was first confirmed in the sun with a spectroscope. Explain how it was possible to use an atomic spectrum to prove the existence of an element.

5 Part B. Flame Tests for Elements Introduction: When atoms in the ground state are heated to high temperatures, some electrons may absorb enough energy to allow them to jump to higher energy levels. These excited state electrons are unstable and they will fall back to their normal positions of lower energy. As the electrons return to the ground state, the energy that was absorbed is emitted in the form of electromagnetic energy. Some of this energy may be in the form of visible light. To do a flame test on a metallic element, the metal is first dissolved in a solution and the solution is then held in the hot, blue flame of a Bunsen burner. This test works well for metal ions, and was perfected by Robert Bunsen. Many metallic ions exhibit characteristic colors when vaporized in the burner flame. The wavelength of a light can be measured in several different units. You may see it expressed in nanometers (nm), meters (m), or Angstroms (Å). These are some important conversion factors: 1.0 nanometer = 1.0x10-9 meters 1.0 Angstrom = 1.0x10-10 meters 1.0 Angstrom = 0.10 nanometers The spectroscope we will be using for this lab measures the wavelength of the bright line spectra in thousands of angstroms. Pre-Lab: Read the lab procedure and introduction, and set-up your lab notebook with the following information Title Objective Data Tables Objectives: 1. To observe the characteristic colors produced by certain metallic ions when vaporized in a flame 2. To identify an unknown metallic ion by means of its flame test. Materials: Potassium chloride Barium Chloride Strontium Chloride Sodium chloride Copper Chloride Calcium Chloride Wooden Splints Safety Goggles Lab Aprons Bunsen Burner Procedure: 1. Put on lab aprons and safety goggles. 2. Fill a beaker halfway with water to depose of used wooden splints. 3. Set up Bunsen burner. 4. Take a wooden split of the element solution out of the labeled beaker and insert into flame. Record the color produced by the element. (Be as detailed as possible) 5. Repeat for all elements. (Obtain the rest of the splints by getting them from the front table) 6. When you have finished testing your eight metallic solutions, obtain a sample of an unknown solution. Perform a flame test and identify the metal present by the color of the flame. 7. Repeat Step 6 for the second unknown and then CLEAN UP!

6 Data: (Copy table in your lab notebook. Remember title, do not cut and paste) Metal solution Color of Flame Potassium Strontium Calcium Barium Sodium Copper Unknown # Unknown # Conclusion: Based on your observations, identify the unknown you examined: Unknown # is Unknown # is 1. Explain the process by which visible light is produced in this lab. Include vocabulary words such as ground state, excited state, quantum, photon, etc. 2. What were your two unknowns? How did you know? 3. Which elements contained high energy electrons? Which elements contained low energy electrons? How could you tell? 4. If milk was boiling on a gas stove and it boiled over, what color would you expect to see? Why? 5. Show your work and remember units and sig figs for the following problems. Planck's constant (h) = x J sec c = 3.00 x 10 8 m/s E = h * υ c = λ * υ a) If the frequency of a red spectrum line is at 1.60 x Hz, what is its wavelength? How much energy does each photon of this light have? b) If the frequency of a violet spectrum line is at 2.50 x Hz, what is it s wavelength in meters? How much energy does each photon of light have? c) What is the frequency of light which has 8.33 x J?