Analyzing Atomic Spectra
Developed by Jenny Joyce, Ferndale High School, Ferndale, WA ~ December 2002

Electromagnetic Radiation and Atomic Spectra

When energized, electrons jump around the nuclei of atoms producing energy in the form of electromagnetic radiation (more commonly called “light”). Electromagnetic radiation comes in many familiar forms; radio, infrared, ultraviolet, X-rays, microwaves, visible light, and gamma rays. Each one of these is unique type of light. If there are an infinite number of electron jumps, a continuous atomic spectrum will be created, producing white light. If there are only a few electron jumps, the atomic spectrum may show only one or two colors of the rainbow. The light produced by electron jumps is called an atomic spectrum.

When atomic spectra were first discovered about one hundred years ago, no one had a clue what they were. One thing, however, was clear. Elements with high atomic numbers (lots of protons and electrons) produced lots of atomic spectra. Elements with low atomic numbers produced simpler atomic spectra.

The Colors of Spectral Lines

Light energy travels in waves, very much like waves in water. Light waves have a velocity, c, (300,000 km/sec in a vacuum), a wavelength, , (lambda, distance from the crest of one wave to the crest of the next wave), and a frequency, f. Wavelengths that are shorter than infrared light are usually measured in nanometers, nm. One nanometer is a billionth (1x10-9) of a meter. The wavelength of the light determines the color, as you can see from the table at the top of the lab write-up.

Hydrogen, for example, the simplest element in the universe, produces only four spectral lines in the visible part of the electromagnetic spectrum. The lines have wavelengths of 410 nm, 432 nm, 486 nm, and 656 nm*.

* It is important to point out here that there are other lines in the electromagnetic spectrum for hydrogen, but these lines lay to either side of the visible light portion (ultraviolet to the right of violet and infrared to the left of red). We can’t see them with our naked eyes.
/ Because each of the 92 naturally occurring elements has a unique number and set of electrons and orbitals, each element produces a unique atomic spectrum. You can therefore identify the elements that are present in a sample by analyzing the atomic spectra produced.
Breaking Light up into Atomic Spectra
A rainbow is formed when raindrops break the sun’s light into the component colors. You undoubtedly have seen how white light can be broken up into a rainbow of colors (a spectrum) with a prism. Light can also be broken into a spectrum by what is called a diffraction grating. A diffraction grating is a piece of transparent glass or plastic ruled with many finely spaced lines (the surface of a CD will do the same thing). If you hold a grating up in front of a light source, the light will be broken up into a spectrum. /

Emission and Absorption Spectra

The spectra that I have described so far, produced by passing energy through any material, are all called emission spectra. Each of the colored lines that you see in an emission spectrum is produced when electrons jump from high orbits to low orbits after being excited. When we view light from any white light source with a diffraction grating, we will see a continuous emission spectrum.

If a cold sample of a gas is placed between the white light source and the diffraction grating, however, a different kind of spectra can be observed. The full spectrum rainbow will be seen to have dark lines spread throughout it. They look like emission lines in reverse. These dark lines in an otherwise continuous spectrum are called absorption lines and they are closely related to the colored lines you see in emission spectra. Atoms absorb light as well as emit light. When white light passes through a cold gas, the gas atoms absorb wavelengths of light that make their electrons jump to other orbits, producing energy gaps, dark lines in the continuous spectrum. The positions of these dark lines correspond exactly to the positions of lines in an emission spectrum of the same gas.

Step One
Full spectrum light waves strike atoms of a cold gas / / / Step Two
The electron absorbs a green wavelength of light and jumps into a higher orbit / / / Step Three
The light continues on minus the wavelengths that were absorbed

Emission and absorption spectra are two tools that are commonly used to identify elements whether they are in Bellingham Bay or on the star Zubenelgenubi. The pictures below show both the emission and absorption spectrum for an element. Notice that the bright emission lines and dark absorption lines are in exactly the same relative position.

FraunhoferLines

Our sun produces a continuous spectrum of light. If you look closely at the spectrum, however, you will see dark absorption lines cutting through the rainbow. These lines were first observed and mapped by a Bavarian optician named Joseph von Fraunhofer and were called (surprisingly) Fraunhofer lines in his honor. The lines indicate that the sun is surrounded by an atmosphere of cooler gases that absorb some of the wavelengths of light coming from the main body. The spectral lines from the sun match the spectral lines seen from elements on Earth, demonstrating that the sun is made from the same stuff as the Earth.

Similar lines are found in the spectra of stars, giving astronomers a powerful tool for studying the chemistry and physical processes that go on outside of our solar system. Spectra produced by stars and galaxies can also be used to figure out how fast objects in the universe move using the Doppler Effect. Below you can see emission line spectra for a smattering of elements.

Analyzing Atomic Spectra Worksheet

Name(s) ______Period_____ Date:______

Questions:

1. Look at the spectrum at the top of this page and determine the colors of the following four spectral lines? These are the four emission lines in the visible part of the spectrum for hydrogen.

410 nm ______/ 432 nm ______/ 486 nm ______/ 656 nm ______

Keep in mind that v =  and E = h

2. Arrange the following bands of light in order from longest wavelength on the left to shortest wavelength on the right.

Radio / Gamma Rays / X-rays / Visible Light / Microwaves / Infrared / Ultraviolet

3. Arrange the same bands of light above in #1 in order from lowest frequency on the left to highest frequency on the right.

4. Arrange the same bands of light above in #1 in order from highest energy on the left to lowest energy on the right.

5. Which color of visible light has…

a) the shortest wavelength? ______c) the least amount of energy? ______

b) the longest wavelength? ______d) the greatest amount of energy? ______

6. The drawing on the right depicts only four of the many possible energy levels around the nucleus of an atom. Answer the following questions:
a) Draw arrows between energy levels for all possible energy transitions for this atom.
b) How many spectral lines will this atom produce from these energy levels? /
c) Which transition (from n = ___ to n = ___ ) will give you the highest frequency (shortest wavelength) of light?
d) Which transition (from n = ___ to n = ___ ) will give you the lowest frequency (longest wavelength) of light?

7. How does a hydrogen atom, which has only one electron, have so many spectral lines?

You will need to complete the following questions online. Go to the following website:

8. When you open this page, you should see the periodic table with the absorption spectrum for hydrogen at the top of the screen.

a) Near the middle of the page you will see two buttons for different kinds of spectra; “absorption” and “emission”. Click back and forth between these two types of spectra. What do you notice about the emission lines compared to the absorption lines?

b) Click on one of the four emission lines in the spectrum for hydrogen and a number will pop up. This is the wavelength (in angstroms) for this particular line. One angstrom is a unit of measurement that is equal to 0.1 nanometers. Therefore, divide the number you see by ten to get the number in nanometers (because there are 10 angstroms in one nanometer). Beginning with the far-left line, write the wavelength (in nanometers) for each line in the space below.

Far left ______Far right

9. Click on sodium and the emission spectrum for this element will appear at the top of the page. When heated, sodium produces a bright yellow light because the two yellow lines you can see dominate its emission spectrum. Sodium vapor lamps are commonly used for bridges and highways for the yellow light they provide. What are the wavelengths (in nanometers) of these two lines?

10. Click on element number 95 on the bottom row of the table. Americium is an artificial, radioactive element used in your home smoke detectors. What color would you predict americium to glow when energized?

11. If you were to heat thorium, element number 90, what color do you think would be produced? Why?

12. Why are spectral lines often referred to as “atomic fingerprints”?

13. Below you will find the known spectra for five common elements followed by the spectrum recorded by a telescope for a distant star. Examine the spectra and answer the questions that follow.

a) How does the light that astronomers see from distant stars and galaxies tell them that the same atoms with the same properties exist throughout the universe?

b) Which element is not in the star that produced the “unknown spectrum”? How can you tell?

14. Check out the two emission spectra below the full spectrum on the right. The top one is of the element lithium measured in a lab. The bottom one is the spectrum of a star taken with a telescope. The two spectra appear similar, but have some obvious differences. What are the differences, and how do you account for them? /
  1. How are the wavelength and the frequency of waves related?

______

  1. In an atom, how is the energy emission (photon release) related to the frequency of the photon that is emitted?

______

  1. What do the different frequencies of light describe? Why is understanding that light is emitted in more than one frequency important?

______

  1. Which color of visible light has the largest wavelength?

______

  1. Which color visible light has the highest energy?

______

  1. What is the atomic emission spectrum, and how is the emission spectrum of an element like a person’s fingerprint? Explain. (see fig. 5.13 in book-page 141)

______

  1. Explain what is meant by ground state and excited state. How do they relate?

______

  1. When color is perceived by a person’s eye, they are actually observing the different frequencies of the light waves. (Respond with either supportive evidence, or an argumentative statement. Give evidence for your claims. Give evidence from the lab AND from either the book or an internet source. – Cite source using endpoints. Place your sources at the bottom of this page. You may write them in pen or pencil.)

______

Sources: