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Santa Monica College
Chemistry 11
Atomic Emission Spectra
Objectives
The objectives of this laboratory are as follows:
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To build and calibrate a simple meter-stick spectroscope that is capable of measuring
wavelengths of visible light.
To use this spectroscope to observe and measure the line spectra emitted by mercury,
hydrogen and other elements.
To use Bohr’s theory to identify the electronic transitions that give rise to each wavelength
observed in the line spectra of hydrogen.
Background
Atomic Emission Spectra
Electrons in atoms normally occupy the lowest energy states possible. Such an atom is said to
be in the ground state. However, electrons can be excited to high energy states when they
absorb energy. This energy can be provided by heat, light, or an electrical discharge. The
electrons will then return to lower energy states, eventually returning all the way to the ground
state. As the electrons return to lower energy states, they release their excess energy. Often,
this energy is released in the form of light, with each atom or molecule releasing a single
photon of light for each electron energy transition it makes.
For example, in the hydrogen discharge tubes used in this experiment the energy of the electric
discharge first dissociates the H2 molecules into H atoms, then excites the electrons in the H
atoms into high energy states. Due to conservation of energy, the amount of energy in an
emitted photon will exactly match the amount of energy lost by the electron as it moves to the
lower energy state.
Different colors of light are associated with different photon energies. For example, a single
photon of blue light has more energy than a single photon of red light. Thus, the color of the
light emitted by a particular atom depends upon how much energy the electron releases as it
moves down to a lower energy level. The energy levels that are allowed for each atom depend
upon the number and arrangement of protons and electrons in the atom. As each element has
different energy states available to it, each element releases photons of different color when its
atoms return to their lower energy states. Since each atom has many excited states (high
energy levels) available to it, several colors of light can be emitted by each element. The set of
individual colors emitted by an element is called its spectrum. Since the spectrum of each
element is unique, spectra can be used like fingerprints to identify unknown elements.
Atomic Emission Spectra
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Santa Monica College
Chemistry 11
Wavelengths of Light
Visible light is only one kind of electromagnetic radiation. The wavelength of radiation
determines what kind of radiation it is. The human eye is able to detect only a narrow range of
wavelengths of electromagnetic radiation, those from about 400 nm to about 700 nm.
Radiation with wavelengths less than 400 nm includes ultraviolet, x-ray, or γ-rays, while
radiation with wavelengths longer than 700 nm includes infrared radiation, microwaves, and
radio waves. In this experiment, we use our eyes to detect the radiation emitted by excited
atoms, and therefore we work only with visible light.
The color of light is related to its wavelength (), its frequency (), and the energy of its photons
(). Shorter wavelengths of light have higher frequencies and higher photon energies while
longer wavelengths of light have lower frequencies and less energy per photon.
It is easy to convert between photon energy, wavelength, and frequency using the relationships:
 = c and photon = h
where c = the speed of light = 2.998 x 108 m/s, and h = Planck’s Constant = 6.626×10-34Jï‚·s.
These two relationships can be combined to give a third relationship:
photon = hc / 
Thus, the spectrum of an element can be described by listing the particular wavelengths of light
that its atoms emit.
To measure wavelengths in a line spectrum we must first separate them. To the naked eye,
the various wavelengths (colors) of light emitted by an element are mixed together and appear
as a single color that is a combination of the component colors. However, if we view the light
through a prism or a diffraction grating, the individual wavelengths are separated. A diffraction
grating is a piece of glass or clear plastic with many very narrow and closely spaced lines on it.
As the light emerges after being reflected by the grating, these tiny lines cause the reflected
light to interfere with itself in such a way that the different wavelengths of the light to appear in
different positions to the left and right of the original direction in which the light was traveling, as
shown in the figure on the next page.
Using a light source that contains known wavelengths of light, we can measure exactly where
each known wavelength appears along a meter stick. Since this position has a linear
dependence upon the wavelength, a graph of wavelength versus position of the spectral lines
will yield a straight line. Once the best-fit straight line has been determined, the equation of this
line can then be used to convert positions of other spectral lines to wavelength. For example,
using the same apparatus, it is possible to view the spectrum of a new element, measure where
Atomic Emission Spectra
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Santa Monica College
Chemistry 11
its spectral lines occur on the meter stick, and then read the calibration graph (or use the
equation of the line) to determine the wavelength to which each of those positions corresponds.
The calibration graph is therefore an integral part of the spectroscope.
Bohr’s Theory and the Balmer Series
For atoms that contain only one electron, the theory proposed by Niels Bohr can be used to
calculate wavelengths for transitions between particular electronic energy levels of the atom. In
this experiment the only one-electron atom we will consider is hydrogen. Applying Bohr’s
theory for hydrogen, a close match can be found between the calculated wavelengths and
those measured experimentally.
To calculate the wavelengths of light emitted by hydrogen atoms, recall that the energy of an
electron in the nth energy level of a one-electron atom is given by:
En = –
Z 2R
n2
where R is the Rydberg Constant (= 2.18 x 10-18 J), Z is the nuclear charge, and allowed values
for n are n = 1, 2, 3,…,∞. For hydrogen, the nuclear charge is +1 so this equation becomes:
En = –
R
n2
The change in energy for the electron when it makes a transition from one level to another is
given by its subtracting its initial energy from its final energy:
Eelectron = Ef – Ei
Atomic Emission Spectra
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Santa Monica College
Chemistry 11
By conservation of energy, the energy of the photon emitted as this electron drops to a lower
level must equal the change in energy for the electron. However, since photon energies must
be a positive quantity, the absolute value of the change in energy for the electron must be used:
Ephoton = | Eelectron |
Once the energy of the photon is known, it is readily converted into a wavelength as discussed
earlier:
photon = hc /  or  = hc /photon
Because there are many energy levels possible for the electron in a hydrogen atom, and
because the electron could jump from any higher n to any lower n, there are many lines in the
spectrum of hydrogen. However, most of these lines occur at wavelengths that our eyes cannot
detect (either infrared or ultraviolet). The visible portion of the spectrum (which you will observe
in this experiment) was the first to be studied by scientists since it is the only portion that can be
seen with the naked eye. This series of spectral lines is named for one of the first scientists to
study it and is called the Balmer series. Note that all of the spectral lines in the Balmer series
involve transitions from a higher n level to the n = 2 level. You will need this information to
complete the calculations for your lab report.
Procedure
Materials and Equipment
Meter sticks bolted in a T shape, diffraction grating, flashlight, dark blanket, 3 ring stands with
ring clamps attached to each, high voltage power supply, hydrogen, mercury, helium and other
polyelectronic element discharge lamps.
Safety
 Exercise extreme caution with the high voltage supplies as severe shocks are possible! Do
not touch the front of the power supply while it is plugged in.
 Do not touch or attempt to remove the discharge tubes from the high voltage supplies. In
addition to the risk of electrical shock, the tubes become very hot with use. The power supplies
must be turned off and unplugged before changing discharge tubes.
 View the light emitted by the discharge tubes through glasses or goggles. Both glass and
plastic lenses will absorb most of the harmful UV radiation emitted by many atoms.
Part A: Constructing and Calibrating a Meter-Stick Spectroscope
Construction of Spectroscope
Work in groups of 4 unless instructed otherwise. Select a workspace on one of the lab
benchtops away from bright light sources.
1. Obtain 3 ring stands and adjust the iron rings so they are all at exactly the same height,
about 5 inches above the bench top.
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Santa Monica College
Chemistry 11
2. Obtain a pair of meter sticks that have been bolted together in a T shape. Place the ring
stands under the ends of the meter sticks so the meter stick arrangement is held about 5
inches above the bench top and is level.
3. Place a high voltage power supply (5000 V – Danger – Do Not Touch If Plugged In!)
containing a mercury discharge tube at the intersection point of the two meter sticks as
shown in the figure below. (Note that this is a very high voltage power supply. You must be
careful never to touch it when it is plugged in. When you need to insert or remove a
discharge tube, turn it off and unplug it before touching the tube. Also note that the tubes
become hot from use. You must let them cool before attempting to remove them.)
4. Mount a diffraction grating held by a rubber stopper in a utility clamp attached to a ring
stand. Place the ring stand so the diffraction grating is centered over the vertical meter
stick and located about 20 cm from the free end of the vertical meter stick (see figure).
5. Be sure not to bump the meter sticks, ring stands, or diffraction grating! If any of these
components is moved during the experiment, the results will be less accurate. You will
need to move the power supply to change discharge tubes, so it is a good idea to mark its’
initial position with masking tape so you can be sure to put in back in the same position
each time.
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Santa Monica College
Chemistry 11
Calibration of Spectroscope using Mercury
6. Be sure you have a mercury (Hg) discharge tube in your power supply and turn it on. When
you look at the lamp through the diffraction grating, you should see four bright lines at
varying positions along the horizontal meter stick: violet, blue, green and yellow. You can
use either the spectrum to the left of the lamp or to the right. It doesn’t matter because they
are the same, but you must be consistent with the side you choose for the rest of the
experiment. Note that the violet line can be hard to see and should be just beyond the blue
line. The wavelengths of these four lines are supplied below.
Color
violet
blue
green
yellow
Wavelength
404.7 nm
435.8 nm
546.1 nm
579.0 nm
7. Measure the position (in cm) of each line observed along the meter stick and record these
positions in the table on your report form.
8. Data Analysis:
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Using Microsoft Excel, plot a graph of the wavelength of the four lines versus their
recorded positions along the meter-stick. Wavelength should be on the y-axis and line
position on the x-axis. Apply a trendline to this data and obtain the equation and R2
value of this line. Staple this graph to your report.
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This equation obtained is your calibration equation. This equation can be used to easily
convert a measured line position (in cm) into a wavelength (in nm).
Part B: The Line Spectrum of Hydrogen – A Single Electron Atom
1. With the power supply unplugged, remove the mercury tube and replace it with a hydrogen
tube. Then plug it in again and turn it on.
2. When you look through the diffraction grating you should see at least three lines clearly:
red, aqua and blue. You may also see a deep violet line (after the blue line). Measure the
position (in cm) of each line observed along the meter stick and record these positions in
the table on your report form. When finished, use your calibration equation to determine the
wavelengths of each of these lines. These are your experimental wavelengths.
3. Data Analysis:
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Using the equations of Bohr’s Theory, calculate the wavelengths of the first six lines in
the Balmer series. These are your theoretical wavelengths.
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Compare your theoretical wavelengths with your experimental wavelengths. Then
identify which electronic transition (n = ? ï‚® n = 2) is responsible for each of the colored
lines you observed in the hydrogen spectrum.
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Calculate your percent error for each line: % error =
Atomic Emission Spectra
| theoretica l  exp erimental |
theoretica l
x 100
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Santa Monica College
Chemistry 11
Part C: The Line Spectrum of a Polyelectronic Element
1. With the power supply turned off and unplugged, replace the hydrogen tube with either a
helium, krypton or argon tube. Write the name of the element you choose on your report
form. The plug in the power supply and turn it on.
2. Look through the diffraction grating and record the color of the 4-6 brightest spectral lines
you see and their positions (in cm) along the meter-stick. When you have finished, use your
calibration equation to determine the wavelengths of the lines you observed. These are
your experimental wavelengths.
3. Data Analysis:
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Go to the NIST government website at
http://physics.nist.gov/PhysRefData/Handbook/element_name.htm.
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Select the name of the element you chose, then click on the box labeled “Strong Lines”.
Scan the wavelength column for the wavelengths you measured to see if you can find any
close matches.
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The tabulated wavelengths are given in units of Angstroms (1 A = 10-10 m). The wavelengths
you have measured in lab are in nm (1 nm = 10-9 m). Thus, you must convert the numbers in
the NIST table to nm by dividing them by 10. (i.e. 5890 A = 589.0 nm)
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Each element actually emits hundreds of wavelengths of light, but only some of those
wavelengths are emitted with enough intensity (brightness) for our eyes to see them. The
tabulated wavelengths include those that we cannot see with the naked eye, so you must
scan down the “Intensity” column in the table of wavelengths to look for those lines that have
high values for intensity. Look for those spectral lines that are significantly more intense than
most since they are the ones you are likely to have seen in lab. Note that this means the
wavelengths may not match perfectly.
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Record the tabulated wavelength for the most intense line nearest to each wavelength you
observed. These tabulated values are your true wavelengths. Then calculate the percent
error for each of your measured experimental wavelengths.
Atomic Emission Spectra
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