Download Spectroscopy of Atoms and Molecules

CHEM 121L
General Chemistry Laboratory
Revision 1.6
Spectroscopy of Atoms and Molecules
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Learn about the Interaction of Photons with Atoms and Molecules.
Learn about the Electronic Structure of Atoms.
Learn about Spectroscopy.
Learn about Beer’s Law.
In this laboratory exercise, we will probe the behavior of electrons within atoms using Emission
and Absorbance Spectroscopy. We will first examine the photons emitted from excited atoms of
various salts. Then we will observe the line spectrum of excited Hydrogen atoms and excited
Helium atoms. Finally, we will leverage the photon absorbance of a solution of Copper Ions to
determine the concentration of those Ions.
Probing the behavior of electrons within atoms is problematic; atoms themselves are far too
small to be seen and their presence must be inferred, and the electron itself is a quantum
mechanical object. However, understanding the behavior of these electrons is important because
this behavior determines an atom’s chemistry. Thus, we must find an indirect probe of the
electronic behavior of an atom. It is found that useful probes of this behavior are the photons of
light which interact with an atom. If an interacting photon’s energy matches that of an electronic
transition within the atom, the photon can be absorbed. Conversely, an electronically excited
atom can relax and emit a photon whose energy matches the atom’s electronic transition. These
photons are directly observable; therefore they provide us with a window on the behavior of
electrons within an atom.
In the realm of molecules, the experiment is much the same; electronic transitions of the
molecule will allow for either absorbance or emission of photons. And, again, the energy of
these photons can be directly measured, giving us insight into the behavior of the electrons in the
molecule.
Photons are quantum mechanical objects that exhibit a Wave-Particle Duality. The wavelength
() of a photon is related to its speed (c) via:
c = 
(Eq. 1)
where  is its frequency. In a vacuum the speed of a photon is c = 2.99792 x 108 m/sec. Its
energy (E) is then related to its frequency via:
E = h
(Eq. 2)
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where h represents Planck’s constant; h = 6.62608 x 10-34 Jsec. Photons can cover a wide range
of energies, from very low energy Radio Waves to high energy Gamma Radiation. This
Electromagnetic Spectrum of photon wavelengths is represented below:
Typical electronic transitions within atoms and molecules are such that the corresponding
photons have energies in the Visible and Ultraviolet regions of the spectrum.
If a photon of the correct Energy impinges on an atom, such that this energy matches the energy
required for an atomic quantum state transition, the photon can be absorbed:
This phenomenon will result in a dark line in the rainbow of colors emerging from a sample
radiated with White Light. Conversely, if an atom that has previously been excited into a higher
quantum state relaxes, a photon whose Energy matches the difference in state energies can be
emitted:
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This will appear as distinctly colored light emitted from the excited sample. In these cases, the
Energy difference between the participating quantum states is related to the Energy of the
photons via:
E = Ephoton = hc / 
(Eq. 3)
Thus, the photons absorbed or emitted by a sample of the atom in question are a direct probe of
the energy difference between quantum states of that atom.
Again, the behavior of electrons within molecules can be probed in pretty much the same
manner.
In the case of emissions from excited atoms, observing emitted photons of multiple wavelengths
implies many quantum states may be involved in producing the spectrum. For example, three
quantum states can produce emitted photons of three different wavelengths:
Visually we see these photons as a characteristic color emitted by the sample. For instance,
excited Sodium (Na) atoms emit visible photons of wavelengths 568.8205nm, 588.9950nm, and
589,5924nm; with the latter two so-called D-Lines being much more intense than the first. Thus,
the yellow color we see emitted by Na atoms is dominated by these two intense D-Lines, which
occur in the Yellow region of the spectrum. (This is why sodium vapor lamps have a yellow hue
to them.)
These photons can be separated by passing the emitted light through a dispersing element, such
as a prism or diffraction grating, to produce a series of spectral lines; each line corresponding to
photons of a given wavelength. Identifying which quantum states are involved in the emission of
photons of a particular wavelength is usually difficult and will not, in general, be considered
here.
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A major complication occurs if an absorbing or emitting species is in a condensed phase; a liquid
or a liquid solution. If the absorbing molecule/atom is in a solution, it is surrounded by
constantly jostling solvent molecules. Thus, each molecule/atom finds itself in a slightly
different environment than its brothers. This causes the energy gap between the quantum states
responsible for the absorbance of photons to be slightly different for each molecule/atom. This
means we will have a series of very, very closely spaced absorbance lines. Practically, this
means the Absorbance Spectrum will be a broad band, rather than a sharp line. This is as
diagramed below:
In this case, we usually identify the absorbance band by the wavelength of maximal absorbance,
max.
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Measuring the wavelength of absorbed or emitted photons is accomplished in a spectrometer.
The spectrometer will also measure the Intensity of the absorbance or emission. This intensity is
then related to the number, or concentration, of the absorbing/emitting species.
A UV-Vis Spectrometer consists of a White Light source, dispersive optics to separate the
wavelengths of the light, a sample compartment and a detector.
Light from the source passes through an Entrance Slit and is focused on the dispersive element,
such as a diffraction grating. This separates out the various wavelengths comprising the White
Light into a rainbow of colors. The desired wavelength is selected by rotating the dispersive
element such that the desired color passes through an Exit Slit, through the Sample and then onto
a detector.
The Transmittance of the light is then defined as:
T = P / Po
(Eq. 4)
This is simply a measure of how many photons pass through the sample without being absorbed.
This is then related to the Absorbance by:
A = - log T
(Eq. 5)
As we have noted, the Absorbance will depend directly on the Concentration (c) of the absorbing
species:
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Additionally, two other factors determine the Absorbance of a sample. The further light must
travel through a sample, the so-called pathlength b, the greater is the Absorbance. The ability of
the absorbing species to absorb light at the given wavelength, known as the extinction coefficient
, a quantum mechanical effect, will also determine the Absorbance. The relationship between
these influences and the Absorbance is given by the Beer-Lambert Law:
A = bc
(Eq. 6)
Note the Absorbance is directly proportional to each factor.
One last point regarding absorbance spectroscopy needs to be addressed; different colored
solutions absorb different wavelengths of light. The color of the light absorbed is directly related
to the color of the light transmitted; i.e., the color we see. The color absorbed is the Complement
of the color we see:
Color Absorbed
Red
Orange
Yellow
Yellowish-Green
Green
Bluish-Green
Greenish-Blue
Blue
Violet
Wavelength (nm)
650-780
595-650
580-595
560-580
500-560
490-500
480-490
435-480
380-435
Color Observed
Bluish-Green
Greenish-Blue
Blue
Violet
Purple
Red
Orange
Yellow
Yellowish-Green
This simply means our spectrometer must be set to a wavelength corresponding to the maximal
absorbance of the absorbing species.
In this lab we will observe the light emitted by atoms that are excited in a hot Bunsen burner
flame. Then we will observe the spectral lines in the emission spectrum of excited Hydrogen
atoms and excited Helium atoms. Finally, we will measure the absorbance of a solution of
Cupric Ions (Cu2+). By knowing how this absorbance depends on concentration, we can
determine the concentration of the solution.
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In our first experiment, we will simply observe the colors of the photons emitted when select
atoms are excited. We will do this by placing a drop of a solution of the salt of the desired atom
into a Bunsen burner flame. The heat of the flame will be sufficient to excite the atoms in
question. We can then observe the light emitted when the atoms relax.
We will be examining emissions from the following atoms:
Li, Na, K, Ca, Sr, Ba, Cu
by testing solutions of the following salts:
LiCl(aq)
NaCl(aq)
KCl(aq)
CaCl2(aq)
Sr(NO3)2(aq)
Ba(NO3)2(aq)
Cu(NO3)2(aq)
Note, in each case we are relying on the fact that neither Cl- or NO3- will influence the emission
from the cations.
In our next experiment we will observe the photons emitted by excited Hydrogen atoms. In this
case we will excite the atoms by placing gaseous Hydrogen (H2) into a discharge tube and
applying a high voltage to the sample. In the resulting electric field, the H2 molecules dissociate
and leave the resulting H atoms in an excited state (H*).
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H2
2 H*
(Eq. 7)
H*
H + photon
(Eq. 8)
These excited atoms will subsequently relax, giving off photons corresponding to the energy
differences between associated quantum states. We will examine the resulting glow by passing
this light through a prism to disperse the photons according to their wavelength.
We will view the resulting spectral lines thru a spectroscope that includes a slit for focusing the
light from the discharge tube, a prism and a rotatable viewer to find the spectral lines.
The first model of the Hydrogen atom to explain these spectral lines was put forth by Neils Bohr
in 1913. In this model, the electron orbits the nucleus in quantized orbits. In its most relaxed or
Ground State, the electron orbits very close to the nucleus. When excited, the electron moves
into a higher energy orbit farther from the nucleus. When it relaxes, the electron moves back to
an orbit closer to the nucleus, emitting a photon in the process.
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Unfortunately this model, although predictive for the Hydrogen atom, has some physically
unappealing features and is not extendable to the electronic behavior of atoms of other elements.
Even Helium, which has only one additional electron, has a much more complex electronic
spectrum. To illustrate this, we will observe the visible lines in Helium’s discharge spectrum.
Finally, in our last experiment, we will measure the absorbance of a solution of Cupric Ions
(Cu2+) using a visible absorbance spectrometer; a Shimadzu UV-2550.
By doing this for a series of solutions of known concentration, we can determine the value of the
Extinction Coefficient, . Subsequently, an absorbance measurement of a solution of unknown
concentration can be used to determine its concentration.
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Pre-Lab Safety Questions
1.
What kind of damage can occur to your eyes due to exposure to Ultraviolet (UV)
Radiation? Below what wavelength radiation should you be concerned about UV
exposure? Above what intensity radiation should you be concerned about UV exposure?
2.
What kind of eye protection must be worn to prevent eye damage due to UV radiation?
3.
Other than the present lab, in which lab were we concerned about eye damage due to UV
radiation exposure?
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Procedure
Flame Tests
1.
Obtain a looped nichrome wire, a Bunsen burner, and the series of salt solutions to be
tested.
LiCl(aq)
NaCl(aq)
KCl(aq)
CaCl2(aq)
Sr(NO3)2(aq)
Ba(NO3)2(aq)
Cu(NO3)2(aq)
Also obtain some 6M HCl for cleaning the nichrome wire.
2.
Light the burner. Dip the nichrome wire into the 6M HCl, then place the loop into the
hottest part of the burner flame. Repeat this step two or three times, until the flame around
the wire is nearly colorless.
3.
Dip the loop of the wire into one of the salt solutions; saving the NaCl solution for last.
Place a drop of the solution into the burner flame. Record your observations. Clean the
wire and repeat the test for each salt solution. (When performing the test on Barium, it may
be necessary to view the flame through a Cobalt Blue filter. This salt is frequently
contaminated with Sodium. Thus, the Sodium emissions must be filtered out.)
4.
Clean the nichrome wire and return it.
Hydrogen Emission Spectrum
1.
Go to the darkened room where your instructor has set-up the Hydrogen Discharge Tube.
2.
Turn on the power supply and observe the emitted light through a spectroscope. You
should see a series of sharp, bright spectral lines. Consult with your instructor to confirm
you are indeed seeing the correct lines.
3.
Record your observations.
4.
View the Helium spectrum through a spectroscope. Record your observations.
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Absorbance Measurements
1.
Obtain an appropriate amount of solid Cupric Sulfate Pentahydrate, CuSO4•5H2O, to
prepare 50mL of a 0.1M solution. You should weigh this out on an analytical balance.
Record the actual mass you use. (Consult with your instructor to confirm your
calculations.)
2.
Use a 50mL volumetric flask to prepare this solution.
3.
Now, fill a spectrometer cuvette. (Handle the cuvette from the edges to avoid getting
your finger prints on the faces. Use a Kimwipe to clean the faces of the cuvettes before
making any readings.)
4.
Now, prepare 50 mL of a 1:2 dilution of your above Cupric Sulfate solution. Fill another
spectrometer cuvette.
5.
Prepare additional 1:2 dilutions until you have a total of 5 filled cuvettes of decreasing
concentration.
6.
Obtain 10 mL of a solution of Cupric Sulfate of unknown concentration. Fill another
cuvette with this solution of unknown concentration.
7.
Proceed to the instrument room to make the absorbance measurements.
8.
Use the Shimadzu UV-2550 scanning UV-VIS spectrometer to measure the absorbance of
each solution at 660nm, the max for the Cupric Sulfate Pentahydrate solution.. Use a
Water filled cuvette as a “blank” to first “zero” the instrument. (Your lab instructor will
assist you in making these measurements. The Shimadzu spectrometer is a scanning,
double-beam instrument. We will be using it to make single wavelength measurements, so
will only be using a fraction of its capabilities.)
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Data Analysis
Flame Tests
No Analysis Needed
Hydrogen Emission Spectrum
1.
Examine the complete emission spectrum for Hydrogen provided in the Appendix.
Determine the wavelengths of the emission lines you observed. (These should all be
Balmer Lines.)
2.
Calculate the photon energies for each spectral line you observed.
3.
Prepare an Energy Level Diagram of quantum states which are responsible for the emission
lines you observed. For this diagram, assume all the Balmer Lines terminate in the same
quantum state. (This state is designated with quantum number n = 2.) Provide quantum
number designations (n = 3, 4, 5, ...) for your quantum states in the Energy Level Diagram.
4.
Identify the orbits in the Bohr Model of the Hydrogen atom responsible for each quantum
state. The Bohr Model provides that the radius of the electron’s orbit is given by:
r = 0.529 x n2 [Angstroms]
(Eq. 10)
where n is the state’s quantum number. Calculate the radius of each of these orbits.
5.
Wavelengths for all the visible spectral lines for Helium can be found at:
http://hyperphysics.phy-astr.gsu.edu/HBASE/quantum/atspect.html
Assign wavelengths to all the spectral lines you observed for Helium.
6.
The following spectral lines are due to the indicated transitions:
Line
587.62 nm
447.18 nm
Transition
3d – 2p
4d – 2p
What is the wavelength of the spectral line corresponding to a 4d – 3d transition? In what
region of the spectrum does this spectral line occur?
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Absorbance Measurements
1.
Using a software package such as Excel, prepare an Absorbance calibration curve using the
Cupric Sulfate solutions of known concentration. To this plot add a Trendline and obtain
the equation of this line.
2.
Using its measured Absorbance and your calibration curve’s Trendline equation, determine
the concentration of the unknown Cupric Sulfate solution.
3.
Determine the Molar Absorptivity  using the Beer-Lambert Law; (Eq. 6).
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Post Lab Questions
1.
Examine the emission spectrum of the Hydrogen atom provided in the Appendix. The first
two Lyman Lines are due to n = 2 to 1 and n = 3 to 1 quantum state transitions. Use the
photon wavelengths of these lines to calculate the energy differences for the associated
states. Do the same for the first Balmer Line, which is due to a n = 3 to 2 quantum state
transition. Are the results internally consistent? Explain. A simple energy Level Diagram
may be useful.
2.
In your spectrophotometer Absorbance measurements, how do you expect the errors in
subsequent measurements to compare to the error in the measurement for the stock
solution?
3.
Identify the Oxidation State for all the metal cations used in the Flame Tests.
4.
What is the color of the Cupric Sulfate solution used in the Absorbance measurements?
What color light is being absorbed by these solutions?
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Appendix - Complete Emission Spectrum of Hydrogen