Analyzing Elements and Substances:
Spectroscopy Skill Building Exercise
Remote Learning
BACKGROUND
In this unit, our focus will be on analyzing elements and substances, and one of the main
questions that we will be investigating involves how we can use visible light, one form of
electromagnetic radiation (EM), to derive information about various substances. More
specifically, you will use visible light to detect, identify, and quantify the amount of an ion or
other substance in various materials. Using analytical methods that are based on the analysis
of different types of EM radiation absorbed or emitted by a substance is called spectroscopy.
EM radiation is a form of energy that can be generated by accelerating particles that have
an electric charge. Different types of EM radiation can be characterized by either wavelength (λ)
or frequency (�). Depending on the wavelength or frequency, EM radiation falls into different
classification, shown in the figure below.
Figure 1. Electromagnetic spectrum.
Chemical substances interact with the various types of EM radiation in different ways.
The type of EM radiation absorbed or emitted by a substance is a differentiating characteristic
of the material, and each type of EM radiation can tell us something different about the
substances. For example, IR spectroscopy uses analysis of EM radiation in the IR region to
provide information about the types of bonds present in a molecule, whereas visible light has
enough energy to excite electrons in a substance. The visible part of the spectrum from 400
(violet region) to 700 (red region) nm will be our focus during this unit.
Absorption and emission of visible light involves particles interacting with light of
specific frequencies. In order for an electron to change its energy level, it must absorb or release
a photon with an energy equal to the difference between two existing energy levels.
Figure 2. Absorption and emission spectra of hydrogen (left), where you can see the
lines have the same wavelength in both, and examples of energy levels (right).
According to the quantum mechanical model of the atom, atoms can exist in different
electronic energy states depending on how their electrons are distributed among various
energy levels. The absorption or emission of energy by an atom is associated with these
electron transitions between different energy levels, and the difference between these energy
levels can be investigated with spectroscopy.
Absorption
The color of most objects depends upon the
interaction between visible light and the electrons of
atoms or molecules that make up the object. It is usually
the result of a dynamic process on the molecular level:
the absorption of light and the resulting change of a
molecule's quantized energy. The object absorbs certain
wavelengths of white light, and we see what is left over.
The particular wavelengths of light that a given
substance absorbs determine the color we perceive in
things like grass, blood and carrots. For absorption, the
perceived color of an object has a complementary
relationship with the color of the visible light absorbed.
For example, a substance that appears green will absorb
light in the red region.
Figure 3. Color wheel with wavelength values
for the visible region.
The fundamental law governing light absorption, known as Beer's law, is: A=εLc where:
A = absorbance of the sample (also defined below in terms of light intensity)
c
=
concentration
of
the
absorbing
species
(mol/liter)
L (or b)= path length traveled by the light beam through the sample (cm),
ε = molar extinction coefficient of the absorbing species (L·mol-1·cm-1)
Beer's law states that the absorbance of light is proportional to the product of the
concentration, c, and the path length, L. The proportionality constant is ε, the molar extinction
coefficient. Normally L (or b) is measured in centimeters and c in moles per liter therefore the
molar extinction coefficient has units of reciprocal centimeters times reciprocal moles per liter.
Beer's Law is useful in the quantitative analysis of solutions containing colored
compounds, like CuSO4 and CoNO3. At a particular wavelength, Beer's Law can be used to
relate absorbance and concentration. Since A = εLc, a plot of absorbance vs. concentration,
for fixed ε and L, will be a straight line with slope=εL and intercept zero. A typical
spectroscopic analysis uses a series of solutions (standards), of known concentrations, to
produce a Beer's Law calibration or standard plot, whose slope can be used to calculate ε.
Then, an unknown sample is analyzed by measuring its absorbance, and Beer's Law is then
used to calculate the concentration of the unknown. These analyses are normally carried out at
the wavelength of an absorption maximum since the absorption coefficient is least sensitive to
changes in wavelength near a maximum.
Emission
Each element has a unique set of emission lines because each element’s electrons have a
unique quantized distribution around the nucleus. The energy and wavelength of the light is
described by the equation E = hc/λ, where λ is the wavelength, h is Planck's constant (6.63 x 1034 J sec), and c is the speed of light (3.00 x 108 m/sec). If you capture the image of light emitted
by an element, the color your eyes see can be broken down into a set of emission lines, and the
emission lines are mathematically described by the equation shown above.
We can calculate the change in energy when an electron is emitted by using the following
equation:
12
12
𝛥𝐸 = −2.178 × 10−18 𝐽 × (𝑍 2 ) × ( 2
− 2
)
𝑛𝑓𝑖𝑛𝑎𝑙 𝑛𝑖𝑛𝑖𝑡𝑖𝑎𝑙
where Z is the atomic number. For hydrogen, Z=1, and the final state is n=2, and the initial states
will be n=3, 4, 5, etc. You should prove this to yourself by reading through your textbook. Also
note that the change in energy will be negative. This is simply showing that energy is emitted
rather than absorbed, so the absolute value of the change in energy should be used for the rest
of the calculations.
Once we know the change in energy, we can use that to predict the wavelength:
Compare these theoretical values to the ones collected experimentally:
You should refer to Zumdahl, pp 545-549 for further reading about the hydrogen spectrum,
particularly figure 12.10.
OBJECTIVES
By preparing for and performing this experiment, you will:
● Use a spectrometer to measure emission intensity and wavelength;
● Describe the relationship between wavelength and color;
● Identify the wavelengths of hydrogen emission lines in the visible spectra;
● Connect atomic theory to observed spectra.
● Create and use a standardization plot;
● Describe the relationship between absorbance and concentration and between
wavelength and color;
● Measure the concentration of ions in an unknown sample.
MATERIALS
Chemicals (hood/balance bench):
Hydrogen Discharge Tubes
Copper (II) Sulfate CuSO4(s)
Cobalt (II) Nitrate Co(NO3)2(s)
Unknown Copper Solution
Unknown Cobalt Solution
Equipment (dispensing room):
Spectrometer Kit (spectrometer, USB cable)
Emissions Fiber Optic Cable
10 mL graduated pipet
Pipet Filler
50 mL Volumetric Flask
You will also need a laptop with LoggerPro software installed. This can be found on the
ELN in the “Logger Pro Downloads” file under the General Information Folder.
PROCEDURES - (Work In Pairs Or 3’s, No 4’s)
EMISSION SPECTROSCOPY:
General Spectrometer Directions (Emission)
1. Check out a spectrometer from the dispensing window. Use the USB cable to connect it
to your computer. Place the fiber optic cable into the cuvette holder.
2. Start the Logger Pro software. Once open, it should automatically detect the
spectrometer.
3. To set up the spectrometer to measure light emission, open the Experiment menu and
select Change Units → Spectrometer: 1 → Intensity.
4. Measure the emission spectrum of a light source.
a. Click
to start data collection.
b. Aim the tip of the cable at the light source.
c. An emission spectrum will be graphed. You may have to adjust the position of
the light source to get a suitable graph to appear.
d. If necessary, move the light source toward or away from the detector so that the
peak emission is less than ~1.5. When you achieve a satisfactory graph, click
. Because distance makes a difference, make sure to estimate and record
the distance between your light source and fiber optic cable tip. Write down
your observations of the emission spectrum in your data table.
e. To store the data, select Store Latest Run from the Experiment menu.
f. When measuring emissions, you should be sure to pay attention to both
wavelength and intensity.
g. To measure the peaks, “Examine” the graphs by selecting Analyze → Examine.
This will give you a slider bar on the screen with a coordinate box.
5. Make sure to save all of your files. Each lab partner should have a copy of the file.
Part 1: Hydrogen Spectrum
NOTE: The power supply that lights the discharge tube generates a LARGE voltage. Avoid
touching any exposed metal connections.
There are one or two hydrogen tubes set up in the lab. These tubes have a short life-span. Please
turn them off immediately after taking your measurements.
Begin by simply looking at the hydrogen discharge tube and recording any observations. Next,
measure the emissions with the spectrometer, adjusting the cable to obtain large peaks. You
should be able to read at least three peaks. There are a number of very weak peaks near 400
nm, so once you measure the large peaks, you may want to adjust the scale of the axes to see
the smaller peaks. Once you have recorded the hydrogen emission lines, you will need to
compare those experimental values to the theoretical ones predicted.
You will need logger pro to open and examine the linked files.
Photos of the Hydrogen tube off and on:
Links to Hydrogen Emission spectra:
Hydrogen File #1
Full spectrum focused on the largest
peak.
Hydrogen File #2
Spectrum focused on second peak
Hydrogen File #3
Spectrum focused on third peak
ABSORPTION SPECTROSCOPY:
General Spectrometer Directions (Absorption)
Calibration
1. Connect the SpectroVis Plus spectrophotometer to the USB port of a computer.
2. Start the Logger Pro software, you will need 3.10 version or newer. Once open, it should
automatically detect the spectrometer.
3. Calibrate the spectrometer.
a. Prepare a blank by filling an empty cuvet ¾ full with distilled water.
b. Open the Experiment menu and select Calibrate → (Spectrometer). The following
message appears in the Calibrate dialog box: “Waiting ... seconds for the device to
warm up.” After 90 seconds, the message changes to: “Warmup complete.”
c. Place the blank in the cuvet holder of the spectrometer. Align the cuvet so that the
clear sides are facing the light source of the spectrometer. Click “Finish
Calibration”, and then click
.
d. You are now ready to collect absorption measurements. More details are below,
but for general measurements: Click
. A full spectrum graph will be
displayed. Note that one area of the graph contains a peak absorbance. Click
to complete the analysis.
Part 2: Concentration of an Unknown Solution
It is your goal in this part to determine the concentration of an unknown solution of copper or
cobalt ions. To complete this task, you will need to:
1. Create a stock solution of your chosen ion.
2. Make dilutions from your stock and measure the absorbance.
3. Plot absorbance vs. concentration using the provided excel file.
4. Measure the absorbance of the unknown corresponding to your chosen ion.
5. Use the slope from your graph to determine the unknown concentration.
For Remote Learning: Use the document below to complete your lab assignment for the
standard curve. You will use cobalt for this experiment.
Link: Cobalt Standard
Creating a Stock Solution
1. Choose either copper or cobalt (TAs may assign an ion to you).
2. Determine the mass needed to make your solution. We have predetermined a good
starting stock. For copper, you will be making 50 mL of a 0.100 M solution. For cobalt,
you will need to make 50 mL of a 0.200 M solution.
3. Measure the mass of your ion, record the value in your notebook, and transfer the solid
to a 50 mL beaker.
4. Add ~25 mL of DI water to the beaker. Place a stir bar in your solution and stir with a
stir plate until the solid is dissolved.
5. Quantitatively transfer your solution to a 50 mL volumetric flask.
a. Pour the solution from the beaker into your volumetric flask.
b. Add ~ 5 ml of DI water to the beaker and swirl around.
c. Pour this small amount into the flask.
d. Repeat (b) and (c) again to ensure all original solution from the beaker is rinsed
and transferred to the flask.
6. Add DI water up to the line on the volumetric flask.
Determining the maximum wavelength (λmax)
1. Empty the blank cuvet and using a plastic transfer pipet, rinse it once with small amounts
of your stock solution, and then fill the cuvet ¾ full with that solution.
2. Align the cuvet so that the clear sides are facing the light source of the spectrometer and
place it in the spectrometer.
3. Click
. A full spectrum graph will be displayed. Note that one area of the graph
contains a peak absorbance. Click
to complete the analysis. You can use command
(mac) or control (PC) E to move your cursor along the graph.
4. Note the absorbance and the maximum wavelength.
5. If the absorbance is greater than 1.0, you will not be able to use the stock as one of your
standards. Your highest concentration should have an absorbance no greater than 1.0,
and the lowest should be around 0.05-0.1.
6. Note: All standards should be measured at the wavelength you just determined.
Diluting the Stock and Preparing the Standards
● Dilutions should be made in large or small test tubes. There are test tube racks in cabinets
against the windows.
● An example of a dilution would be to take 3 mL of your stock and 2 mL of water. A second
dilution could be made with 1 mL of stock and 4 mL of water. You don’t have to make the
total 5 mL, but that is an easy way to approach the dilutions. Using 0.5 mL increments
such as 2.5 mL stock to 2.5 mL of water is acceptable.
● If the initial stock is well above an absorption of 1.0, then you will need to have dilutions
that contain more water than stock.
● Your standards should have absorption values that range from 0.05 to 1.0.
Measuring the Standards
1. Put the cuvet filled with your weakest standard in the spectrometer. Click
When
the absorbance reading stabilizes, record it in your notebook. Note: you should be
reading ALL absorbance values at the λmax that was determined above.
2. Empty the contents of your cuvet into the appropriate waste, and rinse it twice with small
amounts of the next solution to be measured.
3. Repeat the absorbance measurement for each solution.
4. When all standards are measured, plot your concentration vs. absorbance data using
Excel or other plotting program (provided for you this first week). Record the equation
of the line in your in-lab assignment.
5. Make sure to use trial and error with the standards. Solutions should span the
range of 0.05 – 1.0, or as much of it as possible. If there are points that don’t fit,
discard that data and repeat any necessary measurement. You need at LEAST 4
good data points. Good meaning fitting to a linear curve with an R2 valued above
0.98.
6. Now that your standard curve is done, you can measure the absorbance of your
unknown, and use that information to calculate the concentration of ions in it.
Measuring the Unknown
1. Obtain about 5 mL of the unknown solution that
corresponds to the standards you made and measured.
2. Rinse the cuvette twice with the unknown solution and
fill it about ¾ full. Wipe the outside of the cuvette and
place it into the spectrophotometer.
3. Measure the absorbance value of the unknown at the
same wavelength of your standards.
4. Using the slope from your standard curve and Beer’s
Law, determine the concentration of your unknown.
5. Record all values in your notebook.
Remote learning: The absorbance of the unknown solution
was measured to be 0.275.
ASSIGNMENT
Make sure to submit your lab assignment pages before
you leave the lab. This experiment will have questions on the
Unit 2 Summary Assignment.