Chemistry 161
Spectroscopy Lab
Objectives:
1. To study the relationship between wavelength, frequency, energy, and use the Rydberg
equation.
2. To gain experience with some spectroscopic techniques.
Introduction:
Part 1: The electron energy levels in atoms and ions are 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 drop in energy levels for electrons from a higher shell to a lower shell.
Observation 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 it
is valence electrons that are involved in these jumps.
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 individual electrons are in their lowest possible values for "n",
the principal quantum number.
Energy can be added to atoms many different ways. It can be in the form of light, an electric
discharge, or heat. The added energy pushes the electron(s) from a lower energy levels to a higher
(excited) energy levels. The excited electron(s) then emit the energy in the form of
electromagnetic radiation (photons) that we see as light, or feel as heat, as the electron(s) fall back
to lower shells. The light emitted has wavelengths and colors that depend on the amount of
energy originally absorbed by the electrons in the atom.
To reinforce the ideas of energy levels and the mathematical relation between properties of the
light associated with a change in electron energy and the change in the electron’s quantum
number we will use the Rydberg equation. The fact that Bohr could derive the Rydberg equation
from his theory and modify it to make it more general was a major indication of Bohr’s success.
The restrictions on when it can be applied are so severe that it has little practical value now.
However, it is an excellent learning tool, illustrating the principles of the light/matter interaction.
One form of the Rydberg equation expresses the change in electron energy as,
 1
1 
Eelectron   RH  Z 2  2  2 
n

 final ninitial 
where RH = 2.178 x 10-18 J, Z is the atomic number of the element, and the n-values are the
quantum numbers for the initial (higher n value) and final (lower n value) energy levels of an
electron undergoing emission. When an electron gains energy, we call the process excitation.
When it loses energy, we refer to the process as relaxation. The square of the atomic number is
not included in the equation in your text as it restricts its examples to hydrogen with atomic
number equal to one. Not that this equation is only valid for an atom or ion with just one electron.
If the electron energy change was associated with absorption or emission of light, there must be a
conservation of energy.
E phot on  E elect ron  0
Here, a positive change in photon energy means a photon was created while a negative change in
photon energy means a photon of that energy was destroyed. In either case, the photon energy
(absolute value) is related to the photon’s frequency by
E phot on  h  
where h is Planck’s constant with a value of 6.626 x 10-34 Js. The speed of light equation
c   
can be used to relate the frequency of light () to its wavelength (). Use 3.00 x 108 m/s as an
approximation for the speed of light, c.
Part 2: Different solutions have different spectral (color) properties depending on the wavelengths
of light absorbed by the molecules in the solution. For instance in a solution that ‘looks’ red, the
electrons within the molecules are absorbing all wavelengths of light, except red. The red
wavelengths of light are not absorbed, but are being transmitted and pass through so that we ‘see’
red. Absorption or emission of light by a molecule depends on electron transitions between
molecular orbital energy levels, just as absorption or emission of electromagnetic radiation by an
atom is determined by electron transitions between different energy levels in the atom and the
Es for those transitions. Molecular spectra follow rules analogous to the rules for atomic spectra:
energy is absorbed only when the amount of energy provided matches the difference in energy,
E, of 2 energy levels. When an electron goes from a higher to a lower energy state, a photon of
specific wavelength is emitted. Every atom or molecule has a characteristic electronic spectrum
depending on its characteristic Es.
Prelab: Handwrite or cut/paste the procedure into your lab notebook. There are no prelab
questions for this lab.
Procedure for Part 1
1. Using the TV Specs, observe and record atomic emission spectra from H lamp.
2. Sketch the atomic emission spectrum of H. Include color and wavelength.
3. Repeat step 1 with the Hg lamp.
4. Sketch the atomic emission spectrum of Hg. Include color and wavelength.
5. Repeat step 1 with the He lamp.
6. Sketch the atomic emission spectrum of He. Include color and wavelength.
Procedure for Part 2
1. Use a USB cable to connect a Vernier Spectrometer to your computer.
2. Start the Logger Pro program on your computer.
3. Calibrate the spectrometer.
a. Prepare a blank by filling an empty cuvette ¾ full with distilled water.
b. Open the Experiment menu and select Calibrate → (Spectrometer:1). The following
message appears in the Calibrate dialog box: “Waiting … seconds for the device to
warm up.” Wait until the message changes to: “Warmup complete.”
c. Place the blank in the cuvette holder of the spectrometer. Align the cuvette so that
the clear sides are facing the light source of the spectrometer. Click “Finish
Calibration”, and then click
.
4. Conduct a full spectrum analysis of a colored solution.
a. Fill a cuvette ¾ full with the colored solution and place it in the spectrometer. Align
the cuvette so that the clear sides are facing the light source of the spectrometer.
b. Click
. A full spectrum graph of the colored solution will be displayed.
c. Once signal has stabilized, click
d. Examine the graph, noting the peak or peaks of very high absorbance or other
distinguishing features. Compare the color of the solution with the wavelength
associated with that color. Use the ‘rainbow’ chart drawn above.
e. To Save your graph, select Store Latest Run from the Experiment menu.
5. Repeat Step 6 with the remaining colored solutions. Remember to keep a copy of each
graph. To export your data open the file menu and select Export As → CSV (Excel,
InspireData, etc.)… You will want to store these runs onto a USB drive or email them to
yourselves. You will want to store these runs onto a USB drive or email them to yourselves.
Lab Data:
Part 1:
1. Calculate the frequency and energy per photon of the shortest wavelength blue line of
mercury.
2. Calculate the frequency and energy per photon of the lowest energy line of helium.
3. Calculate the frequency and energy per photon of the highest energy line in hydrogen.
Part 2
1. Generate an excel graph of Absorbance (y-axis) vs Wavelength (x-axis) for each solution
(you may place all colored solutions on the same graph if you wish). Clearly label
everything and identify the wavelength of maximum absorbance for each solution.
2. Create a table listing the wavelength of peak absorbance for each sample, the color that
corresponds to the wavelength of peak absorbance, and the color of the sample as you
‘see’ it.
Post-lab Questions and Calculations: Clearly report your results, any graphs, tables, and answers
to your postlab questions. Unless using excel, everything is handwritten. Include a statement of
any problems/errors encountered, and a statement of what you learned from this lab.
1. Briefly explain why the gas discharge tube glows when the high voltage is applied. Why
does the gas not glow when the voltage is turned off.
2. When you look directly at the gas discharge tube through the grating, do you see the color
of a single wavelength of light or a blend of wavelengths.
3. Other than our blending of colors of different wavelengths emitted from a single source,
what is another limitation of using our eyes to detect electromagnetic radiation? (Hint –
look at the electromagnetic spectrum in your text or lab poster)
4. Explain how the emission phenomenon can be used to find out what elements are present
in an unknown sample.
5. What is a continuous spectra? What type of source creates a continuous spectra?
6. Explain why atomic emission spectra are not continuous.
7. How many lines are there in the Hg line spectra? How many did you see? What might be
the reason that these two answers are different?
8. Why do atoms absorb and emit light of the same frequency? In other words, why don’t
atoms absorb light at one set of frequencies and emit light at another set of frequencies?
9. Why does the atomic emission spectrum vary from element to element? Explain how
atomic structure influences the spectrum.
10. In part 2, does the wavelength of maximum absorbance (ie: the color that corresponds to
the wavelength) for each solution correspond to the color of the solution that you ‘see’?
Describe why or why not? (For instance, for the red solution, was the highest absorbance
measured at the red wavelength and describe why or why not?)
Postlab Calculations:
1. Use the Rydberg equation for the following calculations.
a. Make a rough sketch of the electron energy level diagram for hydrogen. Make a
horizontal line for each of the energy levels corresponding to n = 1 through n = 6.
b. In your energy level sketch, make a vertical arrow between the levels that illustrates
an electron transition from n = 5 to n = 2. Questions below, c, d, and e refer to the
same transition.
c. Find the energy change, in joules, for a hydrogen electron undergoing this
transition.
d. Convert the energy change in question b to units of kilojoules per mole of electrons.
e. Find the energy, frequency, and wavelength of the photon emitted from a single
hydrogen atom whose electron undergoes the relaxation. Label each result clearly.
2. The specific wavelengths of light emitted when hydrogen atom electrons relax from high
energy levels (n = 3 to n = ) to the lower energy level of n = 2 are members of the Balmer
series. Show step-by-step calculations for the following questions.
a. What is the shortest possible wavelength in the Balmer (visible) series?
b. What is the longest possible wavelength in the Balmer (visible) series?
3. Find the energy required to ionize a ground state hydrogen atom (ie: energy to remove one
ground state electron from an H atom, or ionization energy). That is, what is the energy
required to make the transition from n = 1 to n = ?
4. Calculate the frequency of light emitted when the electron in an He + ion relaxes from the
third energy level to the first energy level.
5. If an electron in a hydrogen atom relaxes to the second energy level, emitting a photon of
wavelength 410.2 nm, what higher energy level did it start at?