AT Chemistry
Atomic Spectra and Atomic Structure
Introduction
Our present understanding of atomic structure has come from studies of the
properties of light or radiant energy and emission and absorption spectra. Radiant energy
is characterized by two variables: its wavelength, , and its frequency, . The
wavelength and frequency are related to one another by the relation:
c = 
where c is the speed of light, 3.00 X 108 m/sec; wavelength is usually expressed in nm
(10-9 m); and frequency is given in cycles per second (sec-1) or hertz (Hz). From this
equation we can see that wavelength and frequency are inversely proportional to each
other. The figure below illustrates this inverse relationship across the electromagnetic
spectrum:
Radiation of different wavelengths affects matter differently. Infrared radiation
may cause a “heat burn”; visible and near-ultraviolet light, a sunburn or suntan; and X
rays, tissue damage or even cancer.
A particular source of radiant energy may emit a single wavelength, as in the light
from a laser, or many different wavelengths, as in the radiations from an incandescent
light bulb or a star.
Radiation composed of a single wavelength is termed
monochromatic. When the radiation from a source such as the sun, or other stars, is
separated into its components, a spectrum is produced. This separation of radiations of
differing wavelengths can be achieved by passing the radiation through a prism. Each
component of the polychromatic radiation is bent to a different extent by the prism, as
shown in the figure below. This rainbow of colors, containing light of all wavelengths, is
called a continuous spectrum. The most familiar example of a continuous spectrum is
the rainbow, produced by the dispersion of sunlight by raindrops or mist.
It was found that by placing a narrow slit between the source of radiation and the
prism, the quality of the spectrum was improved and the monochromatic components
were more sharply resolved. An instrument used for studying line spectra is called a
spectroscope. A spectroscope contains the following: a slit for admitting a narrow,
collimated beam of light; a prism for dispersing the light into its components; an eyepiece
for viewing the spectrum; and an illuminated scale against which the spectrum may be
viewed.
Most substances will emit light energy if heated to a high enough temperature.
For example, a fireplace poker will glow red if left in the fireplace flame for several
minutes. Similarly, neon gas will emit bright red light when excited with sufficient high
electrical voltage. When energy is absorbed by a substance, electrons within the atoms of
the substance may be excited to positions of higher potential energy, farther from the
nucleus. When these electrons return to their normal, or ground-state, position, energy
will be emitted. Usually, some of the energy emitted occurs in the visible region of the
electromagnetic spectrum, whose wavelength range is about 400-700 nm.
Excited atoms do not emit a continuous spectrum, but rather emit radiation at only
certain discrete, well defined, fixed wavelengths. For example, if you have ever spilled
table salt (NaCl) into a flame, you have seen the characteristic yellow emission of excited
sodium atoms. If the light emitted by the atoms of a particular element is viewed in a
spectroscope, only certain bright-colored lines are seen in the spectrum. The figure
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below illustrates the emission spectrum of hydrogen and is called a bright-line
spectrum.
The observation that a given excited atom emits radiation at only certain fixed
wavelengths indicates that the atom can undergo energy changes only of certain fixed,
definite amounts. An atom does not emit continuous radiation but rather emits energy
corresponding to definite regular changes in the energies of its component electrons. The
experimental demonstration of bright-line atomic emission spectra implied a regular,
fixed electronic microstructure for the atom and led to the Bohr model for the hydrogen
atom.
This experiment consists of four parts. In part I you will calibrate your
spectroscope; in part II you will observe emission spectra of hydrogen and calculate the
electronic transition for each of the hydrogen lines, and mathematically and graphically
verify the Rydberg constant; in Part III you will observe line spectra from flame tests;
and in Part IV you will observe the absorption spectra of several transition metal ions.
Part I.
Emission Spectra
In this part of the experiment, you will use a Vernier Spectrometer (V-SPEC) to measure
the light emitted by selected sources. These sources can be, but are not limited to,
discharge tubes, LEDs, lamps, or luminescent or fluorescent liquid solutions.
The electrons of atoms and molecules exist in specific energy states. The energy emitted
by the excitation of electrons is limited to differences between these states, thus specific
energies of light are emitted. The color of a glowing LED, for example, is determined by
the energy of the emitted light. 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 × 10-34 J
sec), and c is the speed of light (3.00 × 108 m/sec). If you are measuring the emission
spectrum of a gas trapped in a discharge tube, only certain wavelengths of light are
emitted by the gas and the “pattern” that is produced is unique for that substance.
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Objectives
In this experiment, you will


Practice measuring the emission spectrum of a source of light.
Compare and contrast the spectra of various light sources.
Materials
Vernier Spectrometer, w/o light source/ light sources:
cuvette holder attachment LEDs
computer discharge tubes
(optional: fiber optic cable) lamp or flashlight
Procedure
1. Use a USB cable to connect a Vernier Spectrometer to your computer. Make sure
that the light source/cuvette holder has been detached from the spectrometer.
(Optional) Connect a fiber optic cable to the threaded detector housing of the
spectrometer.
2. Start the Logger Pro 3.4.6 program on your computer.
3. In your data table, record the type of light source that you will be testing.
4. Prepare the spectrometer to measure light emission.
a. Open the Experiment menu and select Connect Interface → Spectrometer →
Scan for Spectrometers.
b. Open the Experiment menu and select Change Units → Spectrometer: 1 →
Intensity.
5. Measure the emission spectrum of a light source.
a. If you are not using a fiber optic cable, place the light source near (within a
few cm) the detector opening on the spectrometer. The detector opening is a
round, stainless steel piece that sits in a recessed area of the spectrometer.
b. If you are using a fiber optic cable, aim the tip of the cable at the light source.
c. Click . 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 .
Write down your observations of the emission spectrum in your data table.
6. To store the data, select Store Latest Run from the Experiment menu.
7. Repeat Step 5 with another light source, as directed by your instructor.
8. (Optional) Save and/or print a copy of your test results. Select Save As… from
the File menu and save your experiment file.
9. Select Exit from the File menu to close down Logger Pro 3.4.6.
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Data Table
Trial
Light Source
Peaks or Unique Features
Data Analysis
1. Describe, in detail, the emission spectrum of each light source. Emphasize the
features of each spectrum that distinguishes it from the other light sources that
you tested.
2. Identify the wavelengths of every peak in the graph of each light source.
3. Speculate about how the emission spectrum below might have been produced.
Part III Emission Spectra for Some Metal Salts
There is a short list of elements that, when their electrons are given some extra quanta of
energy, the energy can be “seen” by the naked human eye. A common, visually
interesting, activity in the chemistry lab is to demonstrate this phenomenon by placing
one of these elements in a lab burner flame. A bright color will flash momentarily,
signaling the energy exchange happening in the energy levels of the element. This
activity, for one fairly obvious reason, is known as a flame test. It can be pulled off in
many ways, from drizzling a few crystals of a salt onto a flame to evaporating a drop or
two of salt solution over a flame. Thus, when you heat a drop of sodium chloride solution
in a lab burner flame, for example, energy is added to the electrons of the sodium atom.
This added energy is emitted when the excited electrons in the sodium atoms give off
light and fall back to lower shells. The light emitted possesses wavelengths and colors,
depending on the amount of energy the atoms absorbed. You’d expect that each excited
sodium atom would emit one type of light. But it doesn’t quite work that way because
there are billions and billions of atoms being excited and they produce billions of
excitations and emissions. And not all of the atoms receive the same amount of energy,
thus the excited electrons will travel to different energy levels. This means that the color
produced by the heated NaCl solution is actually a series of discrete emission lines. Each
emission describes one quantum path that an electron can take after it absorbs energy
from the flame. 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 × 10-34 J sec), and c is the speed of light (3.00 × 108 m/sec). If
you capture the image of light emitted by an element heated by a flame, 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.
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In this part of the experiment, you will use a spectrometer to measure the light emission
produced by the salts of several elements as they are heated in a lab burner flame. You
will use this information to draw comparisons between the various substances that you
test.
Objectives
In this experiment, you will
•
•
Measure the emission spectrum of various salt solutions.
Characterize the emission spectra of produced by various elements.
Materials
Vernier Spectrometer, w/o light source/cuvette holder attachment or Vernier SpectroVis
dropper bottles of salt solutions, which may include: LiCl, NaCl, KCl, SrCl2, CuCl2,
BaCl2, CaCl2
computer platinum or nichrome wire flame loop
Logger Pro 3 software
distilled water
fiber optic accessory
dropper bottle of 1 M HCl solution
lab burner
Procedure
1. Use a USB cable to connect a spectrometer to your computer. Connect a fiber
optic cable to the spectrometer.
2. Start Logger Pro 3 (version 3.6 or newer) on your computer.
3. In your data table, record the substances that you will be testing.
4. Prepare the spectrometer to measure light emission.
a. Open the Experiment menu and select Change Units ► Spectrometer: 1 ►
Intensity.
b. Open the Experiment menu again and select Set Up Sensors ► Spectrometer:1.
c. Change the Sample Time to 90 ms. Change the Samples to Average to 1.
d. Click the red box in the upper right hand corner to close the dialog box.
5. Ignite the lab burner and adjust the flame so that it is mostly a blue color. The
flame need not be vigorous.
6. Clean the flame loop by placing one drop of 1 M HCl on the loop itself. Hold the
loop in the burner flame for about ten seconds. Repeat this process twice.
7. Measure the emission spectrum of a salt solution. Note: This step is easier to do
with two people. One person should hold the fiber optic cable and a second
person should work with the flame loop.
a. Click > collect.
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b. Hold the fiber optic cable with your fingers very nearly at the tip of the cable.
Aim the tip of the cable toward the flame, and no close than 5-6 cm from the
flame. If your fingertips are warm, the cable is probably too close to the flame.
c. Select a dropper bottle of salt solution and place one drop of solution on the
flame loop.
d. Hold the loop in the burner flame.
e. If the graph does not show a series of peaks, take the loop out of the flame and
carefully cool it with a few drops of distilled water. Place another drop of the salt
solution on the loop and place the loop back in the flame.
f. When you achieve a satisfactory graph, click stop. Write down your
observations of the emission spectrum in your data table.
8. To store the data, select Store Latest Run from the Experiment menu.
9. Repeat Steps 6 – 8 with the remaining salt solutions.
10. (Optional) Save and/or print a copy of your test results.
Data Table – Flame Test Observations
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Summary
Salt
Metal
Color Flame
Wavelengths of
Emission Lines
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Data Table – Energy Calculations
salt
Wavelength of
emission line (m)
Frequency (1/sec)
Energy (J)
Calculation Guide
a. To convert wavelength, 1 nm = 1 × 10–9 m
b. To calculate the frequency, use the equation: c = λν. c is the speed of light (3 × 108
m/s ),λ is the wavelength in meters, and ν is frequency.
c. To calculate the energy of an emission line, use the equation: E = hν. The value of h,
the , Planck constant, is 6.63 × 10-34 J sec.
Data Analysis
1. Describe, in detail, the emission spectrum of each salt. Emphasize the features of each
spectrum that distinguishes it from the other salts that you tested.
2. Identify the wavelengths of every peak in the graph of each salt.
3. Choose at least one emission line from each salt and calculate its energy. If possible,
use a wide range of wavelengths. Compare the energies of the different emission lines.
Which part of the visible light range possesses the greatest energy?
Part IV. Absorption Spectra
The absorption spectrum of an element is the “mirror image” of its emission
spectrum. The principle is the same. If we took a solution of copper ions, for example,
and passed white light through this solution with a spectroscope on the opposite side of
the solution, we would observe a continuous spectrum interrupted by several dark lines.
These dark lines represent the characteristic wavelengths of light absorbed by the copper
ions (i.e. those that have the exact amount of energy to excite the electrons). These are
the same wavelengths that are released when the excited electrons fall back to ground
state levels producing an emission spectrum. Since copper ions appear green they reflect
the green wavelengths of visible light while absorbing the complimentary color – red.
We would thus expect to see prominent dark lines in the red portion of the absorption
spectrum of copper ions.
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The absorption spectra of six different salt solutions will be demonstrated. Record your
observation in Data Sheet 4.
Data Sheet 4
Fill in the color wheel below:
Name of solution
color
colors absorbed
by solution
(dark lines)
colors passing
through solution
ammonium dichromate
__________
_______________
_______________
copper(II)chloride
__________
_______________
_______________
iron(II)chloride
__________
_______________
_______________
potassium permanganate
__________
_______________
_______________
praseodymium chloride
__________
_______________
_______________
erbium chloride
__________
_______________
_______________
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Questions
1. Why does the flame test for a sodium ion result in a yellow flame, even though green
and orange lines appear in the sodium ion emission spectrum? Explain.
2. If you could observe an absorption spectrum of hydrogen, what would it look like?
Explain.
3. Metal ion solutions also contain anions such as chloride ion. When you observe the
emitted light of metal ion solutions, how do you know the resulting spectra come
from the cation, such as sodium ion, and not the anion, such as chloride ion?
Describe an experiment you could perform to verify your answer.
4. Sodium ion, a common contaminant of potassium salts, often interferes with
potassium ion flame tests. To get around this problem, we view the flame through a
piece of cobalt blue glass to mask the color of the sodium ion flame. Briefly explain
why the blue glass blocks the yellow light of the sodium ion but still allows the violet
light of the potassium ion flame to pass through.
This lab was taken from three sources:
Hall, James F., Experimental Chemistry, 2nd edition, D.C. Heath & Co., 1989.
Metz, Patricia A., “Determining Atomic Emission by Spectroscopy”, 1995, Chemical
Education Resources.
Absorption Spectroscopy kit by Flinn Scientific
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