Laboratory Title:
Your Name:
Concepts Addressed:
Lab Goals:
Lab Objectives:
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Benchmark(s) Addressed:
Materials and Costs:
List the equipment and non-consumable material and estimated cost of each
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Estimated total, one-time, start-up cost: $
List the consumable supplies and estimated cost for presenting to a class of 30 students
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Time:
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Introduction:
http://upload.wikimedia.org/wikipedia/en/8/8a/Electromagnetic-Spectrum.png
http://imagine.gsfc.nasa.gov/docs/teachers/lessons/supernova/supernova_student.html
http://www.pbs.org/newshour/extra/teachers/lessonplans/science/hubble.html
http://en.wikipedia.org/wiki/Introduction_to_quantum_mechanics
http://en.wikipedia.org/wiki/Doppler_effect
Spectroscopy
Quantum mechanics (QM, or quantum theory) is a physical science dealing with the behaviour of matter and energy on the scale of atoms and subatomic particles / waves. QM also forms the basis for the contemporary understanding of how very large objects such as stars and galaxies, and cosmological events such as the Big Bang, can be analyzed and explained. Quantum mechanics is the foundation of several related disciplines including nanotechnology, condensed matter physics, quantum chemistry, structural biology, particle physics, and electronics.
The term "quantum mechanics" was first coined by Max Born in 1924. The acceptance by the general physics community of quantum mechanics is due to its accurate prediction of the physical behaviour of systems, including systems where Newtonian mechanics fails. Even general relativity is limited—in ways quantum mechanics is not—for describing systems at the atomic scale or smaller, at very low or very high energies, or at the lowest temperatures. Through a century of experimentation and applied science, quantum mechanical theory has proven to be very successful and practical.
Another development was the discovery of the Zeeman effect, named after Pieter Zeeman (1865-1943). The Zeeman effect could be interpreted to mean that light waves are originated by electrons vibrating in their orbits, but classical physics could not explain why electrons should not fall out of their orbits and into the nucleus of their atoms, nor could classical physics explain why their orbits would be such as to produce the series of frequencies derived by Balmer’s formula and displayed in the line spectra. Why did the electrons not produce a continuous spectrum?
In 1897 the particle called the electron was discovered. By means of the gold foil experiment physicists discovered that matter is, volume for volume, largely space. Once that was clear, it was hypothesized that negative charge entities called electrons surround positively charged nuclei. So at first, all scientists believed that the atom must be like a miniature solar system. But that simple analogy predicted that electrons would, within about one hundredth of a microsecond, crash into the nucleus of the atom. The great question of the early 20th century was, "Why do electrons normally maintain a stable orbit around the nucleus?"
Niels Bohr Model of Discrete Quanta
In 1913, Niels Bohr removed this substantial problem by applying the idea of discrete (non-continuous) quanta to the orbits of electrons. This account became known as the Bohr model of the atom. Bohr basically theorized that electrons can only inhabit certain orbits around the atom. These orbits could be derived by looking at the spectral lines produced by atoms.
Atomic spectroscopy is an extremely important tool for scientists. Because the electron patterns around every kind of atom are unique, and because these electrons interact with light in different ways because of their different positions, you can determine what kinds of atoms are present in a substance by the kind of light absorbed or emitted by the substance. Every atom has a kind of "fingerprint" in the normal light spectrum that is measured with a device called a Spectrometer. This instrument uses a diffraction grating as a prism, splitting the incoming light into its composite colors.
As an example, a marine ecologist may suspect that the reason many bottom-dwelling organisms are dying in a local harbor is because of a chemical pollutant. She samples the mud and chemically extracts a type of metal ion, but she’s not sure what kind of metal it is. She injects the metal ions into the hot flame of an atomic emission spectrometer and observes two line spectra. The lines correspond to the wavelengths of 563 nanometers (nm) and 615 nm. This combination is the "fingerprint" for tin. The ecologist may then trace the tin to a particular type of ship's paint or a nearby industrial source.
Spectroscopy is vital for astronomers. They can analyze elements in stars and nebulae. Star fuel begins as primarily hydrogen atoms. Under extreme pressure and temperature, these atoms undergo nuclear fusion, forming helium atoms, releasing energy in the form of light. As the star ages, its fuel changes, to include the heavier elements. Helium will be used as a fuel source, fusing into heavier elements, and those into heavier elements. Iron is the heaviest element that can form in stars. Elements heavier than iron are formed in supernovae. What elements are part of the star fuel can be determined by spectroscopy.
In the first part of this lab, students will observe various light sources including tubes that have been filled with various types of gases. As electricity passes through these tubes, the gas glows and light is given off. They will compare the spectra of these gas tubes with incandescent (regular light bulb) sources or fluorescent light fixtures.
In the second part of this lab, students will conduct a flame test of 6 different known salts and one unknown chemical. Based on the flame test analysis, they will be able to identify the unknown chemical both by their spectral signature and the color of the flame.
The Doppler effect, named after Christian Doppler, is the change in frequency and wavelength of a wave as perceived by an observer moving relative to the source of the waves. For waves that propagate in a wave medium, such as sound waves, the velocity of the observer and of the source are relative to the medium in which the waves are transmitted. The total Doppler effect may therefore result from motion of the source, motion of the observer, or motion of the medium. Each of these effects is analyzed separately. For waves that do not require a medium, such as light or gravity in special relativity, only the relative difference in velocity between the observer and the source needs to be considered.
The Doppler effect for electromagnetic waves such as light is of great use in astronomy and results in either a so-called redshift or blueshift. It has been used to measure the speed at which stars and galaxies are approaching or receding from us, that is, the radial velocity. This is used to detect if an apparently single star is, in reality, a close binary and even to measure the rotational speed of stars and galaxies.
The use of the Doppler effect for light in astronomy depends on our knowledge that the spectra of stars are not continuous. They exhibit absorption lines at well defined frequencies that are correlated with the energies required to excite electrons in various elements from one level to another. The Doppler effect is recognizable in the fact that the absorption lines are not always at the frequencies that are obtained from the spectrum of a stationary light source. Since blue light has a higher frequency than red light, the spectral lines of an approaching astronomical light source exhibit a blueshift and those of a receding astronomical light source exhibit a redshift.
The siren on a passing emergency vehicle will start out higher than its stationary pitch, slide down as it passes, and continue lower than its stationary pitch as it recedes from the observer. Astronomer John Dobson explained the effect thus:
"The reason the siren slides is because it doesn't hit you."
Depending on the interest level of your students, you can discuss
· quantum mechanics
· uses of spectrometers in different areas of science from astronomy to molecular analysis
· analytical chemistry of determining the properties of an unknown chemical and comparing them to known chemicals
Materials:
· Spectroscopes or diffraction grating (rainbow) glasses
· Power Source
· Spectrum emission tubes
· Poster of Electromagnetic Spectrum (optional)
· Color pencils
· Student worksheet
· 9 v batteries – 1 per group
· Buzzers – 1 per group
· Masking tape
· String
· Large aluminum pans
· Water
· Small pebbles
· Rubber bands
· Alcohol burners
· Probes
· Goggles
· Denatured alcohol
· Matches
· Wire probes
· Flame Test Chemicals
o Sodium Chloride
o Potassium Chloride
o Lithium Chloride
o Strontium Chloride
o Barium Chloride
o Calcium Chloride
o Hydrochloric acid (HCl)
Procedure
Before the students enter the class –
· Set out color pencils, student worksheet, and spectroscopes (rainbow glasses.
· Plug power source into outlet where all students can easily view it. Set out the spectrum tubes.
Introduction
· uses of spectrometers in different areas of science from astronomy to molecular analysis
· extrapolate red shift and blue shift
Experiment 1
· Instruct the students to find an unobstructed view of the power source.
· Have the students view the light (either incandescent or florescent) and replicate the spectrum with the color pencils on the student worksheet
· Switch off the lights and turn on the power source with one of the spectrum tubes inserted in the power source. Have the students record the color they see without the spectroscope.
· Instruct the students to use the spectroscopes and view the spectral lines. Have the students replicate those lines with their color pencils on the student worksheet.
· Repeat with the other spectrum tubes. Be sure to turn off the power source before removing the first tube and inserting the second tube.
Experiment 2
· Instruct the students to test the buzzer and battery. The buzzer will only work when the negative side of the battery is connected to the negative wire, and positive side of the battery connected to the positive wire.
· Release one of the wires to stop the buzzer.
· Securely tape the buzzer to the battery.
· Securely tape the string to the buzzer/battery
· Go outside.
· Attach the wire to the buzzer to produce sound.
· Add some tape to keep the wires in place.
· Instruct one person in each group to rapidly swing the buzzer over head. The other students in that team listen to the change in sound.
· Rotate though so each student has a change to be the swinger.
Experiment 3
· Give each group an aluminum-roasting pan and fill about 4 inches (~10 cm) water.
· The teams must work together to observe the wave changes. First, instruct the first student to drop a pebble in the water and observe what happens.
· Position 2 students on either side of the pan, and another student to drop the pebble. When the pebble is dropped, instruct the two pan students (one pushing and the other pulling) to move the pan in one direction and observe the waves as they do so. Caution students not to push and pull too quickly because the water will slop out. They may want to practice first.
· Repeat until each student observes that the waves shorten in the direction the pan is moving, and lengthening on the other side.
Experiment 4
· Visit and run demonstration of Doppler Effect with sound:http://www.astro.ubc.ca/%7Escharein/a311/Sim.html#Doppler
Experiment 5
Before the students enter the class –
· Set out goggles, the flame test chemicals, probes, alcohol burners, color pencils, student worksheet, and spectroscopes.
· Wipe the alcohol burners with a wet paper towel (in case of a fuel leak while transporting).
Class – Introduction
· quantum mechanics
· uses of spectrometers in different areas of science from astronomy to molecular analysis
· analytical chemistry of determining the properties of an unknown chemical and comparing them to known chemicals
Class – The experiment
· Instruct the students to find an unobstructed view of the power source.
· Have the students view the light (either incandescent or florescent) and replicate the spectrum with the color pencils on the student worksheet
· Switch off the lights and turn on the power source with one of the spectrum tubes inserted in the power source. Have the students record the color they see without the spectroscope.
· Instruct the students to use the spectroscopes and view the spectral lines. Have the students replicate those lines with their color pencils on the student worksheet.
· Repeat with the other spectrum tubes. Be sure to turn off the power source before removing the first tube and inserting the second tube.
· Instruct the students that they will now conduct the flame test with a partner. Each student will be responsible for testing the chemical in the flame so their partner can view the results with the spectroscope. The partners switch roles. The test may need to be conducted several times before the partners agree on the colors, both the color without the spectroscope and the spectral lines.
· Cover all safety with the students. If any student has hair long enough to put in a ponytail, they need to do so. Hand out rubber bands to those students.
· Put on goggles.
· Strike match and light the alcohol burners (either you light them, or if middle school, they can light the alcohol burners).
· The hottest part of the flame is where the flame is white then blue then yellow then orange then red is the coolest. Put probe in the hottest part of flame for better results and do not touch the wick with your probe.
· Students conduct experiment with the 7 chemicals and record the spectral lines using the color pencils and the color observed without the spectroscope.
· When you and your partner agree on your results for all 7 chemicals, extinguish alcohol burner flame with the thimble.
Discussion and thought questions:
1. What is the color and spectral lines for each of the 7 chemicals? What is the identity of the unknown chemical?
2. Helium was discovered in the Sun's corona during the eclipse of 1868. In 1888, traces of helium were isolated here on Earth. How could scientists determine that this was the same gas that had been identified on the Sun?
3. Compare the results of the various gas tube spectra with the spectrum observed using the standard fluorescent light tube. Based on your results, what gas do you think is used in fluorescent light tubes?
4. Was there any difference between the spectra of the standard fluorescent light tube and the compact fluorescent light fixture? Why do you think this is so?
5. Compare the results of the incandescent light bulb with the spectra of the fluorescent light tube and the compact fluorescent light fixture. Based on the observed spectra, can you think of a reason why the fluorescent lights are considered as more "energy efficient"? Could there be a disadvantage to this?