Lab 10 Electronic Energy Levels Chemistry B1A/Summer 2015

Transcription

1 Lab 10 ic Energy Levels Chemistry B1A/Summer 2015 by Kenward Vaughan and William Daniel Much of what we know about the universe beyond our planet Earth is a result of our understanding of how light, or electromagnetic radiation, interacts with ordinary matter. This study is called spectroscopy. The absorption or emission of light involves a change of energy in matter, and this change of energy involves different wavelengths or different frequencies of light. The different energies give information about the electronic energy levels of electrons in atoms, and changes in vibrational/rotational energies of molecules provides information about bond strengths, bond lengths, and angles. The intensity of light absorbed or emitted gives information about how much of a substance is present. Radio Waves Microwaves Infrared Visible Ultraviolet X-Rays Gamma Rays Frequency/Hz >10 19 Wavelength m m cm to mm mm to µm nm nm m nm - pm <10-12 m <pm Energy Transition in matter Information Provided Nuclear Magnetismproton spin Molecular Structure Rotation of molecules Molecular Structure Vibration of molecules Molecular Structure, Compound identification Energy Levels structure Energy Levels Structure Energy Levels and Nuclear Structure Nucleus Energy Levels Nuclear Structure The energy of light (E) is related to its wavelength (λ, the Greek letter lambda and has length units) and its frequency (ν, the Greek nu with units of cylces/second or hertz (Hz)). c = λ ν E = h ν = h c/λ Here c is the speed of light which is X 10 8 m/s and h is Planck s constant which is X Js. The following table breaks down visible light into different wavelengths. Wavelength / nm Color Wavelength / nm Color 650 Red 475 Blue 590 Orange 445 Indigo 570 Yellow 400 Violet 510 Green In this lab experiment you will observe light emission from several different sources and from different ionic compounds where the light is characteristic of the metal cation. A spectroscope or spectrometer Lab_10_ic_Energy_Levels_v4 Page 1

2 will use a diffraction grating to disperse the emitted light into different wavelengths. Spectroscopes typically disperse the light and display the intensity of light (proportional to the concentration) on the y axis and either wavelength or frequency on the x axis. In this experiment the intensity will not be measured. The light emitted by ionic compounds will be observed with the naked eye. Light sources used in the United States include incandescent lamps and fluorescent lamps. Any object above absolute zero temperature emits electromagnetic radiation. An incandescent lamp uses a hot tungsten filament to emit a continuous spectrum of light. The flow of electrons, called electric current, through a tungsten filament, heat the filament to a high enough temperature so the tungsten emits a broad spectrum of visible light. Halogen lamps use iodine or bromine to extend the life of the tungsten filament, and they run hotter and have more of a blue color due to the higher frequencies or energies emitted by the higher temperature filament. Since common incandescent lamps are inefficient (about 95% of their power is emitted as heat and 5% is visible light), their use should decline due to increasing energy efficiency standards. Fluorescent lights work by an electric current providing energy to excite electrons in atoms to higher energy levels, which then relax or fall down to lower states emitting light. The light given off is characteristic and can be used to identify the material emitting the light. Neon lights refer to lights where electric current ionized different gases, (such as neon, argon, and many others), and they emit different colors. What we call fluorescent lights use mercury which emits light that is absorbed by materials called phosphors that emit different colors. These lights have a higher cost for the bulb but the lifetime cost is much less due to their greater efficiency. Materials Spectroscopes Variable Power Supply H, He, Ar, Hg, Ne, Kr, Gas Tubes Incandescent light bulb and CFB LiCl, BaCl 2, CaCl 2, SrCl 2, Zn(NO 3 ) 2, NaCl, MgCl 2, and one unknown salt Procedure Before lab, look up the Ångstrom (Å) unit and determine how to convert this unit of length to nanometers. Your report asks a question about Ångstrom the scientist. You will move around to different stations to observe light emission from different materials. I. Light bulbs a) Incandescent Lamp. Vary the temperature of the filament in the incandescent light bulb by varying the electric power put through the blub. Start with the center knob of the Lab Volt Power Supply (LVPS) turned completely counter clockwise (CCCW). The bulb should be plugged into the socket that shows V. Observe the bulb with the naked eye and gradually increase the voltage to the bulb by turning the LVPS knob clockwise (CW). Lab_10_ic_Energy_Levels_v4 Page 2

3 Record your observations about the color and intensity of light as power (and temperature) are increased. Repeat the above experiment, observing the light bulb through the spectroscope. When looking through the eyepiece, the slit allowing light to enter is on the opposite left side of the spectroscope eyepiece. The light bulb should be very close, but not touching the slit. Inside the spectroscope you will see on the right side what is called a reticule scale. This scale is used to measure the wavelength of light in Ångstroms The numbers should be from 40 to 70 Ångstroms (Å). Record your observations as you increase the temperature of the bulb with the LVPS. It may be convenient to record bands of color and their wavelengths. b) The compact fluorescent bulb. Again, observe using the naked eye and then the spectroscope, the compact fluorescent bulb as you increase the voltage (potential energy). II. Fluorescent Lights a) Hydrogen emission light. Record the color emitted by the hydrogen light with the naked eye. Use the spectroscope to record the wavelength of light of the three brightest lines to the correct number of significant figures. b) Fluorescent lights. Several different gas lights will be available (He, Ne, Ar, and Hg, and the room fluorescent lights). Record the color observed with the naked eye and record the different colors observed with the spectroscope. III. Light Emission from Salts When substances are heated in a flame, electrons can be promoted from their low energy ground states to excited energy levels or states. The electrons can then relax and as they move from the excited energy levels to the lower energy levels they can emit light energy characteristic of that substance. In this lab, metal cations will be emitting the visible light. In a fume hood there will be watch glasses with different salts and Bunsen burners. Each watch glass should have a Nichrome wire with a loop. To be sure the wire is not contaminated with sodium insert the wire loop into the tip of the inner flame (the hottest part of the flame). As the wire gets hot it will emit light, but if the wire is contaminated with sodium you will see an orange color coming up off of the wire. Heat the wire in the flame until you no longer see this color. After this cleaning put the hot loop into the salt and return the wire to the flame tip. Observe and record the color of the emitted light for each substance. It may also be helpful in later experiments if you record observations such as the complete formula including waters of hydration and the compound s appearance. Observe LiCl, BaCl 2, CaCl 2, SrCl 2, Zn(NO 3 ) 2, NaCl, MgCl 2, and the unknown compound. Lab_10_ic_Energy_Levels_v4 Page 3

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5 Lab 10 ic Energy Levels Name Circle Lab Section MW TRam TRpm TRevening Partner s Name Pre-lab exercise 1. Give Ångstrom s full name, briefly state his contribution to science, and show how to convert 58.4 Å into nanometers. 2. Convert the wavelength of blue and yellow light into frequency. Which color is higher in energy? 3. The general form of the Rydberg equation (below) describes the relationship between the wavelength of light emitted and the two electron energy levels in hydrogen. 1 λ = R h 1 n n 2 2 Where n 2 and n 1 are integers and n 2 is the initial, higher energy level and n 1 is the final, lower energy state, and R h = X 10-2 nm -1. In lab you will observe the Balmer series of emission lines which is part of the hydrogen spectral series where the final energy level has n 1 =2, and the light emitted is visible. Look up the hydrogen spectral series and name two other series for hydrogen emission and answer, a) what type of light these series belong to (e.g. are they visible, infrared, UV, X-Rays) and b) what is n 1 equal to (or what is the final electron state). Lab_10_ic_Energy_Levels_v4 Page 5

6 4. Calculate the expected wavelength (λ) in nanometers of the light associated with an electron moving from the 5 th energy level to the 3 rd energy level, (i.e., n = 5 going to n=3). 5. The relationship between the energy of light and the frequency, and wavelength is E = nhν = nhc/λ Here n is the number of photons, h is Planck's constant, ν is frequency, c is the speed of light, and λ is wavelength. Calculate the energy in kj/mol if a mole of these problem # 4 transitions occurred (hint, since the energy is per mole, how many photon are your using in your calculation). Given ΔH fus for ice water is about 4.0 kj/mol, what mass of ice could be melted by the energy of one mole of this electron transition? 6. The arrows represent transitions of an electron from a higher energy level to a lower energy level, and they represent the emission of energy. One transition corresponds to emission of yellow, one corresponds to blue, and one to orange light. Label the arrows as to the color of light they represent. n = 3 n = 2 n = 1 Lab_10_ic_Energy_Levels_v4 Page 6

7 Results Record your observations in your lab notebook I. a) Incandescent Lamp 1. For the incandescent lamp, what was your naked eye observation when you increased the voltage (potential energy) on the LVPS with regard to intensity, colors and spectrum? 2. What was your observation and conclusions about the spectrum when looking through the spectroscope? I b) The Compact Fluorescent Bulb (CFB) 1. What were your observations from the CFB? Lab_10_ic_Energy_Levels_v4 Page 7

8 2. Based on your observations, does a fluorescent bulb work the same way as an incandescent lamp? Explain. II. Fluorescent Lights a) 1. Hydrogen emission light Record the wavelength in Ångstroms of the 3 brightest hydrogen lines, and convert these wavelengths to energy. Wavelength of line (longest 1 st )/ Å Color of H Line Energy / J * Calculated E/J for 3 2 transition (Assigned to East Bench) Calculated E/J for 4 2 transition (Assigned to middle Bench) Calculated E/J for 5 2 transition (Assigned to West Bench) X X X X X X 2. * Using the experimentally measure longest wavelength, show the experimental energy calculation of the energy. 3. **Show the theoretical energy calculation for the energy of the transition assigned to your bench. See equation page 5 for calculating wavelength. Lab_10_ic_Energy_Levels_v4 Page 8

9 4. Starting with the longest wavelength, use the experimental energies compared to the theoretical energies to assign the three different colors to the transition. For example, state if the experimental energy of the green line matches the calculated energy of the n=3 n=2 transition, or the 4 2, or whichever transition. Longest Wavelength- Medium Wavelength- Shortest Wavelength- 5. Draw a potential energy diagram, similar to the pre-lab problem 6, showing the hydrogen energy levels and label with the n value. Label the 3 different transitions with the color of the light emitted. 6. The n quantum numbers represent different electron shells. As you move down the periodic table, the outer, valence electrons are in shells with greater n numbers, which corresponds to the period number. As you move down the periodic table, metallic character of elements increases. Using your potential energy diagram, speculate on why it is easier to move electrons through larger, more metallic solid elements, (i.e., why does electrical conductivity increase as you move down the periodic table). Lab_10_ic_Energy_Levels_v4 Page 9

10 b) Fluorescent Lights (He, Ne, and Hg) 1. Describe how the number of visible emission spectra lines changes as atoms get larger. 2. Explain why the number of visible emission spectra lines observed changes as atoms get larger. Lab_10_ic_Energy_Levels_v4 Page 10

11 III. Light Emission from Salts Briefly state the colors observed for the different known and unknown compounds heated in the flames. Identify the unknowns based on their emission colors. Unknown A = Lab_10_ic_Energy_Levels_v4 Page 11