What Color of Light Has the Lowest Energy
Lab report for students in Dr. Volland's class should be submitted using the equiz
All rights reserved, Dr. Walt Volland, Revised March 31, 2005
The energy levels in atoms and ions are the key to the production and detection of light. Energy levels or "shells"exist for electrons in atoms and molecules. The colors of dyes and other compounds results from electron jumps between these shells or levels. The colors of fireworks result from jumps of electrons from one shell to another. Observations of light emitted by the elements is also evidence for the existence of shells, subshsells and energy levels. The kinds of light that interact with atoms indicate the energy differences between shells and energy levels in the quantum theory model of the atom. Typically the valence electrons are the ones involved in these jumps.
The "quantum" theory was proposed more than 90 years ago, and has been confirmed by thousands of experiments. Science and education has failed to clearly describe the energy level concept to almost four generations of citizens. This experiment is an exercise aimed at throwing a little more light on the subject. ( Don't laugh too hard at the joke.)
Atoms have two kinds of states; a ground state and an excited state. The ground state is the state in which the electrons in the atom are in their lowest energy levels possible (atoms naturally are in the ground state). This means the electrons have the lowest possible values for "n" the principal quantum number.
Specific quantized amounts of energy are needed to excite an electron in an atom and produce an excited state. The animation shows the opposite of excitation. It shows how the excited hydrogen atom with an electron in the n = 3 shell can release energy. If the electron in hydrogen only drops to the n = 2 shell the energy matches a pulse of red light.
Note the size of the electron cloud in the excited atom changes when the electron moves from shell to shell. The size of the atom decreases in volume when the electron goes from the n=3 shell to the n = 2 shell. On average the electrons are closer to the nucleus for lower values of "n". The electron cloud is related to the most probable distance between the nucleus and the electron. The most probable distance increases with increasing "n" value. The excited electron is still "in" the atom even in an excited state. The valence electron will only escape the atom if the electron is given an amount of energy equal to the ionization energy for that atom. (Be sure to view Animation) return to top of page
Energy can be added to atoms many different ways. It can be in the form of light, an electric discharge or heat. This added or extra energy is emitted when the excited electrons in the atoms give off light and fall back to lower shells. The light emitted has wavelengths and colors that depend on the amount of energy originally absorbed by the atoms. Usually each individual excited atom will emit one type of light. Since we have billions and billions of atoms we get billions of excitations and emissions. return to top of page
Not all atoms in a sample will absorb or be excited exactly the same. For example in hydrogen the ground state has the electron in the n= 1 shell. The electron in some hydrogen atoms may be excited into the n = 2 level. Other hydrogen atoms can have the electron excited into the n = 4 shell.
Different elements emit different emission spectra when they are excited because each type of element has a unique energy shell or energy level system. Each element has a different set of emission colors because they have different energy level spacings. We will see the emission spectra or pattern of wavelengths (atomic spectra) emitted by six different elements in this lab. We will then identify an unknown element by comparing the color of the unknown with the flame color of our knowns.
You need to know that white light is the combination of all colors of the spectrum.
Each color has a characteristic wavelength. The wavelength is the distance between the beginning and end of a complete cycle of the light wave. All colors of light travel at the same speed, 3.0 x 108 meters/ second. The animation shows how a prism separates photons of red light from photons of blue light. The photons of different colors fall in different positions on the color spectrum. The position is determined by the wavelength. (Be sure to view Animation)return to top of page
return to top of page
Arizona State University has a nice page on Light at this URL.
http://acept.la.asu.edu/PiN/rdg/readings.shtml
Colorado State University physics department has an animation of the H atom at this URL
http://www.colorado.edu/physics/2000/quantumzone/lines2.html
Definitions of terms and properties of light are at this site return to top of page
http://www.aspsky.org/html/tnl/35/atom.html
The speed of light frequency and wavelength are related by the formula
c = l n
where c is the speed of light 3.00 x 108 meters/ second.
n is the frequency of the light wave in cycles per second
Lambda, l, is the wavelength in meters/cycle.
The frequency, n , equals the number of cycles that pass an observer in a second. These individual cycles each carry a specific small amount of energy. The particles of light are called photons. The energy of a light photon or particle is different for each color. If you have ever seen a Star Trek movie you may have heard of "photon torpedoes" . The science fiction writers suggest that bundles of light energy could be harnessed and used as projectiles.return to top of page
We can calculate the energy of a photon or wave packet using the Planck energy relation shown here. Planck won a Nobel Prize for this work. This relation allows us to tell how much energy a single photon can carry.
| E = h n |
| here E is the energy of the photon in joules |
| h is Planck's constant 6.624 x 10-34 joule seconds |
| n is the frequency of the particular light wave |
return to top of page
Red light has longer wavelength and is lower in energy than blue light. The wavelength of red light corresponds to the range of 700 to 600 nanometers, (7000 Ångstrom or 0.0000007 meters).
Blue light has shorter wavelength in the range of 400 nm (4000 Ångstrom or 0.00000004 meter,
1 Å = 1 x 10 -10 m = 0.0000000001 meter = 1 x 10-1 nanometer). return to top of page
S pectroscopy is the analysis of light spectra and the way in which light interacts with matter. When light is analyzed it is commonly separated into its component colors. The light source is directed on a slit and the "beam" of light is separated using a prism or grating.
The reason that the images are lines is that the light from the lamp is focused on a narrow slit. The illustration shows the separation of a light beam into its component colors. return to top of page
This produces an image of the slit which has the shape of a line. The resulting beam of light can be broken into the color spectrum. or into its components of the spectrum emitted by the atom. You can see the specific colors emitted by the light source. A white light source will give a spectrum like the one shown above. return to top of page
The emission spectrum for sodium shows only two colors in the visible color range. The two colors are yellow and have wavelengths of about 590 nm. The continuous spectrum and calibration scale is shown to give approximate wavelength values.
return to top of page
A CD-ROM mirrored surface behaves like a grating or prism. If you look at the surface of the CD-ROM under a light, you can see the color spectrum.
Street lights are sodium vapor lamps in many communities. These lamps have an orange yellow tint. You can see from the emission spectrum why the sodium vapor lamps would appear yellow and not white. These lamps consume less energy than the older blue colored mercury vapor lamps. Mercury vapor lamps have been sold in hardware stores for yard lighting.
One of the odd things that we sometimes notice is that colored things have a different appearance in natural daylight than they do under mercury vapor or sodium vapor lamps. This is reasonable because the daylight includes all of the wavelengths of white light and the vapor lights only emit a few specific colors that can be reflected into our eye off of any illuminated article.
Example calculations: return to top of page
Frequency
| How to calculate frequency, n , from wavelength and the speed of light, c . |
| A wave with the color red has a wavelength of 6.50 x 10-7 meters / cycle), what is the frequency of the wave? |
| Use c = l n solve for n and we get the following n = {c / l } |
n = {c / l } = (3.00 x 108 meters/ second ) / ( 6.50 x 10-7 meters/ cycle)
n = 4.62 x 1014 cycles / second with three sf
The energy of a photon or single pulse of light energy can be calculated using the Planck energy relation
E = h n = ( 6.624 x 10-34 joules/ cycle) ( 4.62 x 1014 cycles / second )
E = 3.05 x 10-19 joules / cycle
return to top of page
Note : A common household 100 watt light bulb uses 100 joules every second. Even though this energy is not converted completely into light, we can make a rough assumption that it is. If this energy is emitted as light ( really much is disipated as heat) it means 3.28 x 1020 photons (330,000,000,000,000,000,000 photons) would have to be emitted every second to transfer the 100 joules in a second. The numbers are huge. By way of comparison t he population of the wor;ld is only 8,000,000,000 (eight billion) . The atoms and their electrons are constantly changing energy. The electrons are going through ceaseless energy level jumps.
In the second part of the experiment you will observe the color of light emitted by excited gases of elements in sealed glass tubes called "spectrum" tubes. Direct current, DC, high voltage electrons are used to excite the atoms in the spectrum tube. High voltage means 1000 to 2000 volts. This is more than 10 times normal household voltage which is 120 volts AC.
The excited atoms release the energy they gained. Some of this energy is in the form of heat and some is in the form of light. The billions of excited atoms release energy. Each excited atom releases a single pulse of light energy as it returns to the "ground" state or low energy state. There are so many pulses emitted the light appears to be continous.
The excited atoms do not all emit the same energy light because the amount of energy that excited them may differ, but there are limitations on the colors they do emit. The kind of light depends on the size of the gaps between the "shells" or energy levels in the atom. The electrons are changing "n" values in the atom. Remember "n" can have only positive whole number values like 1, 2, 3, ... up to infinity.
The kind of light energy that can be emitted by excited atoms is unique for an element. The pattern of "lines' or colors emitted can be used to identify an element. An powerful extension of this is the ability to measure amounts of an element by measuring the brightness of the emitted light.
A spectroscope can separate the light produced by an emission tube. The color seen by the naked eye is a combination of a number of colors of light. These are separated by a prism or a diffraction grating which acts like a prism. The emission lines can be seen when you look through the spectroscope at the light source. You will be able to observe the "line" spectrum for the elements and record the spectral lines.
How do these emission spectra compare in terms of colors and numbers of emission line positions? Are the spectra identical? What if anything is similar? What is different?
Examine the spectra for the elements Na, Ne, Hg or He and answer the following questions. FILL IN THE FOLLOWING TABLE WITH YOUR ANSWERS Element with greatest number of visible emission lines __________________ Longest wavelength in the spectrum of this atom in nanometers. __________________ Color of light for this longest wavelength __________________
Examine the spectra for the elements Na, Ne, Hg or He and answer the following questions. FILL IN THE FOLLOWING TABLE WITH YOUR ANSWERS Element with fewest number of emission lines __________________ Longest wavelength in the spectrum of this atom in nanometers __________________ Color of light for this longest wavelength __________________ What suggestions do you have additions or changes to this experiment? What "new" idea did you learn from this experiment? Why does a sodium vapor street light look yellow instead of white? What would you expect to happen to the size (volume) of a hydrogen atom when the outer electron moves from the n = 2 shell to the shell with n = 4 ? Would the volume increase?
Revised March 31, 2005 All rights reserved. Dr. Walt Volland (Copyright 1999-2005)
Dr. Walt Volland (Copyright 1999-2002)
What Color of Light Has the Lowest Energy
Source: http://www.800mainstreet.com/spect/emission-flame-exp.html