Astronomy Outreach at UT Austin
EMISSION SPECTROSCOPY LAB
Introduction:
According to the Bohr atomic model, electrons orbit the nucleus in fixed paths with specific energies. Each path is therefore often referred to as an energy level. 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 and become excited. 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. Sometimes the excited electrons acquire sufficient energy to make several energy level transitions. When these electrons return to their ground state, several distinct energy emissions occur. Electrons may become excited when a sample of matter is heated or subjected to an electrical current. The energy electrons emit when returning to the ground state is electromagnetic radiation, or EMR. Sometimes the EMR can be seen as visible light.
In 1900, Max Planck studied visible emissions from hot, glowing solids. He proposed that light was emitted in packets 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 (ν) of emitted light and Planck’s constant (h).
E=h•v
When white light passes through a prism its component wavelengths are bent at different angles. This produces a rainbow of 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. Each band represents a specific energy level change of electrons in the atoms. Since the atoms of each element contain unique arrangements of electrons with unique energy levels, these bands form characteristic patterns, specific to each element called bright line or emission spectra.
In this experiment you will use a spectroscope to examine several continuous and bright line spectra.
Objectives:
1. Explore the relationship between atomic structure and emission spectroscopy.
2. Use a spectroscope and observe sources of continuous and bright line spectra.
3. Calculate the wavelength, frequency, and energy of some of the spectral lines for the elements tested.
Pre-lab: Answer in your own words.
1. According to Bohr’s atomic model, where are electrons found in an atom?
2. What does it mean for electrons to become ‘excited’? List two methods of doing this.
3. State the equation that is used to determine the energy content of a packet of light of specific frequency. Use words, then repeat with symbols. If only the wavelength of the light is known, what additional equation is needed?
4. What form of energy emission accompanies the return of excited electrons to the ground state?
Equipment
Project STAR spectrometer spectrum tubes fluorescent light source
Incandescent light source scientific calculator
CAUTION: 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.
Procedure:
1. As your instructor energizes each tube or light bulb, look at the light with your spectrometer. Your eye will go at the smaller end, and you will point the wider end at the light with the light source in the slit.
2. While looking at the spectrum of colors in the spectrometer, sketch what you see in data table 1, under “Spectrum Pattern”, being careful to represent the numbers associated with the position of each bright line.
3. In data table 2, record the colors and scale readings of at least four different colored lines from the sources in data table 1.
4. Calculate the energy of the spectrum lines of the different colors of light that you chose. A sample calculation follows. Perform work clearly on a separate sheet and summarize in data table 2 below:
a. Change the reading on the spectroscope scale to a wavelength in meters The numbers on the scale range from 350 to 700 nm (10-9 m)
ex. Scale reading is 620: wavelength (() = 6.20 x 10-7m
b. For each wavelength, determine the frequency of the line. Use c = 3.00x108 m/s
c (speed of light) = ( (wavelength) • ν (frequency) or: ν = c/(
ex. ν = [pic]
c. For each wavelength, determine the energy of the line. This represents the energy difference between the excited state of the electrons and the level which they “fell” down to. Use h = 6.63 x 10-34 J•s
E (energy) =h (Planck’s constant) • ν
ex. E = (6.63 x 10-34 J•s) (4.84 x 1014 s-1) = 3.2 X 10-19 J
Data and calculations
Data Table 1.
|Light Source |Spectrum Pattern |Color of Plasma |Element |
|Incandescent Bulb | | | |
| | | | |
|Fluorescent Bulb | | | |
| | | | |
|Tube 1 | | | |
| | | | |
|Tube 2 | | | |
| | | | |
|Tube 3 | | | |
| | | | |
|Tube 4 | | | |
| | | | |
Data Table 2. Perform work on another sheet and attach.
|Source |Color of line |Wavelength(m) |Frequency (Hz) |Energy (J) |
| | | | | |
| | | | | |
| | | | | |
| | | | | |
| | | | | |
| | | | | |
| | | | | |
Post-lab QUESTIONS: Answer in your own words. Be sure to explain all answers.
1. What is the difference between a continuous spectrum and a bright line spectrum?
2. What causes spectral lines? Answer on the atomic level and use Planck’s and Bohr’s ideas.
3. Why do elements have more than one spectral line? Why aren’t there infinitely many lines?
4. In this lab you saw three spectral lines for H, about seven for He, and many for Ne. What does this suggest about the electrons or the energy levels in these atoms?
5. Before helium was discovered on Earth it was found in the sun. How do you think it was found and why do you think it was found on the sun and not the Earth first?
6. You cannot see the electromagnetic radiation that reaches your TV but the wavelength is about 5.0 m. Calculate the energy associated with this EMR.
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