Visible Spectroscopy
Objective : Investigate the emission spectra or pattern of wavelengths (atomic spectra) emitted
by two different elements, study the emission and absorption spectra of colored solutions, and
explore an application of absorption spectroscopy using calibration curves and Beer’s Law.
Introduction
Part 1: 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. 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 the valence electrons are the ones 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 electrons have the 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. 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.
Part 2: Absorption or emission of ultraviolet or visible 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 definite wavelength and frequency is emitted. Every
atom or molecule has a characteristic electronic spectrum depending on its characteristic Es.
Part 3: As seen in Part 2, different solutions have different spectral properties. In this portion
of the experiment those properties will be utilized to determine the concentration of an unknown
sample. In looking at the absorbance of a solution it can be noted that the absorbance of a
solution is directly proportional to the concentration of the solution and the length of solution the
light has to pass through. This relationship is known as Beer’s Law and can be expressed as
follows:
A = εlc
Where A is absorbance, ε is the molar absorptivity (in L/mol*cm), l is the path length (in cm), and c
is the concentration (in mol/L).
A calibration curve can be generated by measuring the absorbance of a series of standard
solutions, and plotting Absorbance (Abs) versus concentration, with Abs on the y-axis and
concentration on the x-axis. The linear regression for this plot will yield a familiar equation of the
line y = mx + b. In this case, we can relate this to the equation above and think of it as A = εl * c +
b. Then by measuring the absorbance of an unknown solution, we can plug it into our equation and
solve for the concentration of our unknown.
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Procedure :
Part 1
1. Using the TV Specs, we will be observing and recording atomic emission spectra from He
and Hg lamp.
 Draw the theoretical emission spectra of He. Include color and wavelength information.

Draw the theoretical emission spectra of He. Include color and wavelength information.
2. As the data for these steps, note and record which lines are actually observed for He and
Hg directly on the spectra drawn above. The spectrum might have more lines than observed.
Use the closest values.
Part 2
The primary objective of this part is to compare and contrast the spectra of various
colored solutions. We will measure and analyze the visible light absorbance spectra of
various samples of aqueous dye mixtures to determine the absorbance spectrum for each
sample.

Sketch the visible spectra (“the rainbow”) here. Include color and wavelength
information.
3. Use a USB cable to connect a Vernier Spectrometer to your computer.
4. Start the Logger Pro program on your computer.
5. 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
.
6. 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.
7. Repeat Step 6 with the remaining colored solutions. Remember to keep a copy of each graph.
Copy your files onto Jackie’s USB (find the correct folder!) and download from the class
website. Make sure everyone in your group knows what the file names are.
Part 3
The primary objective of this part is to determine the concentration of an unknown cobalt (II)
chloride, CoCl2, solution using a calibration curve. You will first measure the absorbance of a
standard solution over the visible light spectrum and select the wavelength of maximum
absorbance (λ max). You will prepare five cobalt (II) chloride solutions of known concentration
(standard solutions) and measure the absorbance of the standard solutions at the λ max. The graph
of absorbance vs. concentration for the standard solutions will describe a direct relationship,
known as Beer’s law. Finally, you will determine the concentration of an unknown CoCl2 solution by
measuring its absorbance with a spectrometer and using the best-fit line equation of the Beer’s law
curve to calculate the unknown’s concentration.
8. Obtain ~40 mL of 0.20 M CoCl2 solution and distilled water in separate beakers.
*If you had to make this solution using solid CoCl2, describe how you would prepare 500 mL
of 0.20 M CoCl2 solution.
9. Label five clean, dry, test tubes 1–6. Use pipets to prepare four standard solutions
according to the chart below. Thoroughly mix each solution with a stirring rod. Clean and
dry the stirring rod between uses.
Test Tube
number
0.20 M CoCl2
(mL)
Distilled H2O
(mL)
1
10
0
2
8
2
3
6
4
4
4
6
5
2
8
10. In Test Tube 6, obtain ~ 5 mL of an unknown sample.
Concentration
(M)
0.20
0.16
0.12
0.08
0.04
11. Use a USB cable to connect a Vernier Spectrometer or a Vernier SpectroVis to a computer.
12. Start the Logger Pro 3 program on your computer.
13. 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.” After the allotted time, the message changes to: “Warmup complete.”
c. Place the blank in the spectrometer; make sure to align the cuvette so that the clear
sides are facing the light source of the spectrometer. Click “Finish Calibration”, and
then click
.
14. Determine the  max for CoCl2 (aq) and set up the data collection mode.
a. Rinse and fill a cuvette ¾ full with the 0.20 M CoCl2 solution and place it in
the spectrometer.
b. Click
. A full spectrum graph of the CoCl2 solution will be displayed. Note that
one area of the graph contains a peak absorbance ( max). Click
to complete
the analysis.
c. To save your graph of absorbance vs. wavelength, select Store Latest Run from the
Experiment menu.
d. Click the Configure Spectrometer icon, , on the toolbar. A dialog box will appear.
e. Select Abs vs. Concentration under Set Collection Mode. The  max will be
automatically selected. Click
to proceed.
15. Collect absorbance-concentration data for the five standard solutions.
a. Leave the cuvette in the spectrometer. Click
. When the absorbance reading
stabilizes, click
. Enter “0.20” (the solution concentration). Click
.
b. Using the solution in Test Tube 2, rinse and fill the cuvette ¾ full. Wipe the cuvette
and place it in the spectrometer. When the absorbance reading stabilizes, click
. Enter “0.16” and click
.
c. Repeat these steps for the remaining test tubes of the standard CoCl2 solution. And
your unknown solution. When you have finished testing the standard solutions, click
.
16. To export your data open the file menu and select Export As → CSV (Excel, InspireData,
etc.)… Copy your files onto Jackie’s USB (find the correct folder!) and download from the
class website. Make sure everyone in your group knows what the file names are.
___________________________CUT HERE ____________________________
Pre-lab Questions:
PL1) Sketch the following quantum leaps using up/down arrows. Then predict which of the
following transitions produces the longest wavelength spectral line:
n = 2 to n = 1, n = 3 to n = 2,
n = 4 to n = 3
PL2) Calculate the energy assosicated with n = 5 to n = 1.
PL3) Calculate the wavelength associated with n = 4 to n = 6
PL4)The yellow light given off by a sodium lamp used for public lighting has a wavelength of 589
nm.
a. What is the frequency of this radiation?
b. Calculate the energy of one photon of this radiation.
c. Calculate the energy produced from one mole of sodium
PL5) Using the sample data below, make a plot of Absorbance vs [CoCl2] and determine the
equation (y=mx+b) for the calibration curve.
Test Tube
[CoCl2] (mol/L)
Absorbance
1
0.20
1.50
2
0.16
1.25
3
0.12
1.00
4
0.08
.75
5
0.04
.50
PL6) Using the calcibration curve from PL5) what is the concentration of [CoCl2] if your unknown
solution has absorbance of 0.88? Show this on the calibration curve and mark it (PL6).
PL7) Using the calcibration curve from PL5) what is the concentration of [CoCl2] if your unknown
solution has absorbance of 3.66? Explain how you got the answer.
Analysis
Part 1
A1) Calculate the frequency and energy per photon associated with the blue line of mercury.
A2) Calculate the frequency and energy per photon associated with lowest energy line of helium.
Part 2
For each colored solution,
A3) Using Excel, calculate Absorbance (A=log 100% / %T) for each of %T readings.
A4 )Using Excel, draw %T (y-axis) vs. wavelength (x-axis). Clearly label everything and identify
the wavelength of minimum transmittance.
A5) Using Excel, draw Absorbance (y-axis) vs. wavelength (x-axis). Clearly label everything and
identify the wavelength of maximum absorbance.
Part 3
A6) Create a graph showing the data and linear-regression equation for the standard solutions.
Make sure to include y=mx+b somehow…
A7) Calculate the molar concentration of your unknown CoCl2 solution, include an explanation of
your calculation of the unknown’s concentration.
Post lab questions
(Q1) What is the relationship (directly, inversely, or not related) between wavelength and
frequency?
(Q2) What is the relationship (directly, inversely, or not related) between wavelength and energy?
(Q3) Do you see the color that is at the wavelength most absorbed or most transmitted? i.e. When
something appears to be blue, is it because the blue wavelength is absorbed the most or
transmitted the most?
(Q4) In your own words, what is the purpose of calibration curve? What information do you need
to construct a calibration curve?
Summary
Part 1
ONE data table including the colors and wavelengths of all theoretical lines and experimental lines.
Part 2
Site the page numbers for graphs. Make sure everything is properly labeled on the graph.
ONE data table including color, the significant wavelengths and absorbance values.
Part 3
Site the page number for the calibration curve. Make sure it includes y=mx+b.