Light & Spectra Lab
Theory
Astronomy is nearly entirely a study of light. From this single source of
information, we can often find the physical conditions, compositions and
processes in distant objects. To do this, we use spectrographs, which allow us to
see how much light an object is emitting at each wavelength. The pattern of light
an object emits at each wavelength is called its spectrum. By spreading out light
into different wavelengths, we can attempt to figure out what is controlling how
much light we see at each wavelength.
To try and figure this out, we look for three components of a spectrum:
continuum emission (or blackbody radiation), emission lines, and absorption lines.
The important thing to know about absorption and emission lines is that every
atom of a particular element (hydrogen, say) has a unique pattern of lines. The
spacing of the lines is the same in both absorption and emission; the only
difference is emission lines are added to the continuum, while absorption lines
are subtracted from it. In contrast, continuum emission always has a smooth
profile, with emission at a given wavelength being only a bit stronger or weaker
than its neighbors.
In this lab, you will be using spectroscope to look at various elements.
These spectroscopes have the same effect as prisms: light of different
wavelengths can be bent different amounts when they pass through the
spectroscope. When you look through the spectroscope, you will see the object
you are looking at right in front of you. As well, on either side of the object you
should see images of the object in different colors. The colors should be ordered
like a rainbow, from blue on one side to red on the other. The brightness of the
image in each color shows you how much light the object is emitting of that
color.
We'll be using these spectroscopes to look at "spectral tubes". These are
tubes that contain gases composed of different elements. By plugging these tubes
into the wall, we can send electricity through them, which adds energy to the
gas. This causes the atoms to become "excited", which is how we describe atoms
whose electrons have been raised into high energy levels. After a short amount of
time the electrons drop back into lower energy states, releasing a photon to carry
off the extra energy. Depending on the number of transitions in each atom and
the energy levels in it, photons of different wavelengths and thus different colors
are released from each gas. The patterns in the types of photons each gas emits
will be the main topic of this lab.
We will also be looking at the colors created by the burning of various
elements. These elements will be in salt solution form. When we burn a sample
of each solution, the burner flame will change color. This color change is due to
exciting the atoms by the addition of the thermal energy. Just as in the spectral
tubes, the electrons will become “excited” and raise to a higher energy level.
Then, after a short period of time, the electrons will fall back to its lower energy
level. When this occurs, a photon will carry off the extra energy and we see the
color change.
Objectives
Students examine emission from spectral tubes and a continuum source through
diffraction gratings or glasses and use their experience with the spectral tubes to
visualize the different energy levels of atoms. Students will also examine the
color change of burner flames due to the burning of various elements in solution.
Procedure
Pre-Laboratory Exercise
Using a compact disc, observe the emission of the following sources of light:
Filament Light bulb
Sun
Street Lamp
Flash Light
Bic® Lighter
Record your observations in your laboratory notebook.
Part I - Spectral Tubes
1. Once you have received a spectroscope and become comfortable with how to
use it, the teacher will darken the room and turn on a number of "spectral tubes"
located at various points around the room, as well as identifying a source of
continuum radiation for you to examine.
2. Look at all the spectral tubes, as well as the continuum source, and draw a
bright-line spectrum for each source. Be certain to label each spectrum that you
create. Things to think about as you draw your spectrums from each source:
Are the lines closely packed, or spread out over many different colors?
Are there many lines you can see, or only a few?
How do the colors of the lines from each tube relate to the color you see
from each tube when you don't look through the gratings?
3. Locate the unknown spectral tubes. Draw a bright-line spectrum for each.
Part II - Flame Test
1. Dip the tip of the wire into the salt solution. Place the tip of the wire into the
center of the Bunsen burner. Observe the color emitted. Repeat for all available
salts solutions.
2. Locate the unknown salt solutions. Dip the tip of the wire into the unknown
salt solution. Record your observations.
**Pre-Lab: Go to WebAssign and complete assignment
Data/Observations
Pre-Laboratory Exercise—write down general observations in your lab
notebook
Part I—create a bright-line spectrum for each element in your lab
notebook as demonstrated below:
RED
ORANGE YELLOW GREEN
BLUE
INDIGO
VIOLET
Part II—write down general observations in your lab notebook
Questions from Part I
1. What is the difference between the spectrum you see from the spectral tubes
and the spectrum you see from the continuum source?
What causes the two spectra to look so different? Describe the physical processes
that lead to the creation of continuum or emission line spectra.
2. The patterns we've seen in emission spectra in this lab and in the absorption
spectra drawn below are controlled by the spacings of energy levels in atoms.
Since photons are absorbed and emitted when electrons move between these
energy levels, the spacing between those levels determines the kinds of photons
that element can emit. Answer the questions below about an atom with energy
levels like those shown to the right. (note: we've labeled each level with a
number n, which gets bigger as the energy of the level gets larger.)
Draw arrows on the diagram to the
right showing all the transitions an
electron could make which would
result in the emission of a photon.
How many emission lines could we
observe from this atom? __________
Which transition would emit light with
the highest frequency (shortest
wavelength)?
From n = _______ to n = ________ .
Which transition would emit light with
the lowest frequency (longest
wavelength)?
From n = _______ to n = ________ .
3. Examine the following spectra:
Solar Spectrum from Sun (Emission spectrum)
Laboratory Spectum of Iron (Absorption spectrum)
What is the evidence for the claim that iron exists in a relatively cool outer layer
of the Sun?
4. Spectral lines are sometimes referred to as "atomic fingerprints". What is it
about each element that causes them to have unique spectra?
5. Promotional signs are often made by bending colorfully glowing tubes into
words or pictures. We usually call these "neon signs". However, you've seen in
this lab that a tube with pure neon emits a very distinctive spectrum, and
appears orange to the eye. How do you think a sign with colors other than
orange in it is made?
6. Consider the energy level diagram that let you predict which energies (lines)
existed for hydrogen. Which energy level transitions did you see? Which energy
level transitions did you feel? Which energy level transitions would do damage
to your skin if your sat in front of the hydrogen tube long enough?
7. Identify the unknown spectra tubes. Be certain to provide the reasoning for
your decision.
Questions from Part II
1. List all the ions that were tested from the Alkali metals.
2. Now list them with their emitted colors from lowest energy to highest energy.
3. How does this compare to their positions on the Periodic Table?
4. Repeat steps 1-3 for the Alkaline Earth metals.
5. Explain the relationship between the energies of light emitted and the
positions of the elements on the periodic table. In your answer refer to the
principle quantum number of the very outermost electrons (valence electrons) of
the elements.
6. Paper logs soaked in solutions of metal salts and dried are sold for fireplaces.
Which ions would give the best colors that would be most visible for your added
enjoyment?
7. Identify the unknown elements. Be certain to provide the reasoning for your
decision.
Sources
Adapted from the Astronomy Department at the University of Washington