Lab # _____
Spectra of the Atoms
Introduction:
The normal electron configuration of atoms or ions of an element is known as the “ground state”. In this most stable
energy state, all electrons are in the lowest energy levels available. When atoms/ions in the ground state are heated to
high temperatures, some electrons may absorb enough energy to allow them to jump to higher energy levels. The atom
is then said to be in the “excited state”. This excited configuration is unstable, and as the electrons return to their
normal (ground state) levels, the energy that was absorbed is emitted (released) in the form of electromagnetic
radiation. Some of this energy may be in the form of visible light.
The color of this light can be used as a means of identifying the elements involved. Such crude analysis is known as a
flame test. Only metal atoms, with their loosely held electrons, can be excited in the flame of a lab burner. Thus, flame
tests are useful in the identification of some metallic ions. If two metals are present in the same solution, the color of
one flame may obscure that of the other. If cobalt glass plates are used, it is sometimes possible to absorb one color
and not the other. In this exercise a characteristic color of a metallic ion will be observed. This color is a mixture of light
of different but specific wavelengths. A more sophisticated method of analysis is to separate this light into its individual
colors (wavelengths) using a spectroscope.
Equipment:
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Spectroscope
Spectrum tubes
Lab burner
wood splints
Cobalt glass
Materials:
0.1 M solutions of:
 Barium chloride
 Calcium chloride
 Copper (II) sulfate
 Lithium chloride
 Potassium chloride
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Sodium chloride
Strontium chloride
unknowns
Safety:
Wear safety goggles throughout the lab. Be aware of the open flame and the high temperature of the Bunsen burner
during and after use. Avoid touching any of the solutions in the lab. Do not allow the wood splints to burn.
Part 1: Flame Tests of Representative Elements
PRE-LAB EXERCISE
Study Table 1 on the following page and then use it to predict the color of the flame produced when each of your test
solutions is heated in a Bunsen burner. Make a prediction for each metal ion listed below Table 1:
Table 1.
Flame spectra of the alkali and alkaline-earth elements.
Predictions: Copy the table below and include it in your lab
Metal ion
Sodium
Predicted flame color
(Na+1)
Barium (Ba+2)
Calcium (Ca+2)
Cesium (Cs+1)
Potassium (K+1)
Strontium (Sr+2)
Lithium (Li+1)
Magnesium (Mg+2)
Rubidium (Rb+1)
notebook. Be sure to give it a relevant title.
Procedure for flame tests:
Flame test observations:
Copy the table below and include it in your lab notebook. Be sure to give it a relevant title.
Metallic ion
Sodium (Na+1)
Potassium (K+1)
Lithium (Li+1)
Calcium (Ca+2)
Strontium (Sr+2)
Barium (Ba+2)
Copper (Cu+2)
Unknowns:
A
B
C
Color in flame
ID of unknown
Procedure for flame tests:
A.
B.
C.
D.
E.
F.
G.
Use a piece of scrap paper to make a diagram of the location of each of the known and unknown solutions
(there are a total of twelve) in the well plate.
Take your well plate to the stock station and fill four of the well plate depressions approximately ¼ full with
four different known solutions. Return to your lab station, taking care not to allow the solutions to “slosh”
and possibly contaminate one another.
Dip your nichrome wire into the acid solution at your lab station and heat it in the burner flame until it is
glowing red to eliminate any possible contaminates.
Dip the loop end of the wire into one of the known solutions and expose the solution to the flame.
Immediately observe and record the color of the flame just above the wire loop. Repeat if necessary to
make accurate and thorough observations.
Repeat steps c. through e. for the remaining known solutions before returning to the stock station. Follow
the same procedure for each known solution. Test the copper solution last as it is most likely to
contaminate the nichrome wire.
Obtain the three unknown solutions as you did for known solutions. Each unknown has only one metal ion
present. Perform a flame test and identify the unknown metallic ion based upon your observations of
known metallic ions. You may want to do direct comparisons by observing the flame of the unknown next
to the flame of a known metallic ion after you have narrowed down the possibilities for the identity.
Clean Up:
Rinse your well plate containing excess solutions into the waste container provided. Wipe out your plastic tray
with a wet paper towel and wipe down the lab counter. Return all materials to the stock station.
Part 2: Spectroscopic Observations of Gases
A spectroscope separates the emitted (released) light into its parts (wavelengths) producing a bright line spectrum
unique for each element. Each bright line represents the energy difference (quantum) as an electron falls from the
exited state to a lower energy state.
For example, in the hydrogen atom:
Electron transition
Color seen
Increasing energy
in n levels
3 to 2
Red line

4 to 2
Blue line

5 to 2
Violet line

6 to 2
Violet line

The simplest type of spectroscope is called a diffraction grating. Like a prism, a diffraction grating
bends light of different wavelengths to different degrees so that the light is separated into its
component wavelengths. When a gas is excited by an electrical current, the bright-line spectrum can
be observed through a diffraction grating.
Use colored pencils or pens to record your observations of the bright line spectra of hydrogen gas,
helium gas, and neon gas in the table below. Cut out the table and tape it into your lab notebook.
Include a title for the diagram.
H2
He
Ne
Wavelength
Color
(nm)
400
violet
500
blue
green
yellow
600
orange
700
red
Increasing wavelength ()
Decreasing energy
Summary:
Answer the questions below in complete sentences in your lab notebook.
Relate your observations of the bright line spectra to Bohr’s model of the atom. Specifically:
1. Why are there distinct bright lines as opposed to a continuous spectrum?
2. What is happening to the valence electrons to cause the bright line spectrum?
3. Why does helium have a bright line spectrum different from the bright line spectrum of hydrogen?
4. How does the bright line spectrum for hydrogen support the Bohr model of the atom?