Lab 6
Name_____________________________
Spectroscopy
Pre-Lab Assignment
This written pre-lab is worth 15% of your lab report grade and must be turned in to your lab instructor before class begins.
1.
2.
3.
4.
Read the entire lab handout carefully. This material will NOT be covered very extensively in lecture, so it will be a selfstudy. Since it is more difficult, set aside plenty of time to read and study this handout.
Make a list of questions on concepts that need to be clarified before lab.
List the suspected safety hazards in today’s lab AND the safety precautions that should be taken to protect yourself from
these hazards. Put this list in a table: one column for safety hazards, one for related safety precautions.
Write a brief summary of the experiment that will be performed today.
Experimental Question
How can the wavelength and energy of emission lines from various elements be determined from the
known wavelengths of a hydrogen spectrum?
Learning goals
•
•
•
Describe how a line spectrum is obtained from gaseous elements.
Describe how the Bohr model explains the presence of lines and gaps in a line spectrum.
Calculate the energy of the line from its wavelength and explain how energy of the line relates to its color.
Background
You have all experienced neon signs and rainbows. You have seen the names of businesses in gas tubes of different colors
and we are used to using the term “neon lights.” Neon is only one of the many gases that are used in these tubes. It is the one
responsible for the tubes that have a reddish orange color. The other colored “neon tubes” actually use other gases to create
the color. A rainbow is created when the light from the sun is broken down into its characteristic components by a prism
effect of small raindrops. In today’s lab, you will be looking at tubes filled with a gas and using a prism to see what colors are
really involved in the color that we see with the naked eye.
You have been learning (or will learn) about atomic structure and you have learned that different elements contain a different
number of protons, neutrons and electrons. The number of protons determines which element we are looking at, and in the
neutral atom the number of electrons is equal to the number of protons. Chemical reactions are concerned with electrons.
Since our eventual goal is to understand chemical reactions, we must first understand the electrons in atoms that are involved
in chemical reactions. In today’s experiment, we will look at atoms (not reactions) in which the electrons are “excited” to a
higher energy state than normal and result in our seeing energy in the form of colored light.
Neon lights are orange, sodium street lamps are yellow and halogen lamps are bright white. Electricity causes the electrons of
these elements to produce light with a distinctive color by causing the electrons to “jump” to orbits of higher energy.
According to the Bohr model, excited electrons emit energy in the form of light when they fall back to lower energy orbits.
Light travels in waves much like those seen on the surface of the ocean before they crash onto the shore. The distance from
wave peak to wave peak is called the wavelength, (λ). You've probably seen ocean waves with wavelengths of 3 meters or
more. Visible light is commonly expressed in wavelengths of 300 to 700 nanometers. These wavelengths are exceedingly
small, but still more than a thousand times longer than the diameter of the atom from which they originate.
Figure 1: Wavelength of electromagnetic waves.
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Lab 6: Spectroscopy
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When light is passed through a prism (or diffraction grating) the light is separated into its component colors. Each color
represents a packet of energy that has a specific wavelength. The colors that are in the visible region of the electromagnetic
spectrum are always in the order of the ROY G. BIV (First letters of Red, Orange, Yellow, Green, Blue, Indigo, Violet) just
like a rainbow. When all the colors are present with no breaks we call it a continuous spectrum (full rainbow) which we see in
sunlight. A rainbow in the sky appears when the sun is at the correct angle (57 degrees) and water droplets are of the right
size then we see the refraction of light through and the reflection of light off the backs of the rain drops. Did you notice that
the sun is always behind you when you see a rainbow? Scientists found from the late 1800’s on that every element has a
unique line emission spectrum when heated. What we mean by a line spectrum is that the light emitted from a source as only
a few set wavelengths (colors) and not an entire rainbow spectrum like the sun.
400nm
700nm
A.
Green
Yellow
Orange
Red
red
Blue
teal
B.
Indigo
violet
Violet
Figure 2. A continuous spectrum has no breaks in it just like a rainbow. (shown in A.). A line emission spectrum
only shows specific lines of color. (shown in B.). The lines shown above are from a hydrogen gas discharge tube.
Notice that this spectrum is ROY G.BIV backwards.
A diffraction grating is used to separate the visible light into its separate colors. When viewing a light source through the
diffraction grating one can see the separated colors that make up the light we see. In spectroscopy the wavelengths of visible
light are recorded in nanometers (nm, x10−9m or a billionth of a meter) . Visible light goes from about 400 nm (violet) to 700
nm (red). Another unit used often in atomic dimensions as well as in wavelengths is the angstrom (Å, x10−10 m) which gives
visible light from 4000 Å to 7000 Å. 110
Most light sources are created by heating elements to a high temperature. Examples are incandescent light bulbs (passing
electrons through a resisting wire), gas discharge tubes (passing electrons through a gas), stars (nuclear fusion of hydrogen
and gravitational collapse) and burning fuels (gas lanterns, candles). Fluorescent bulbs are a type of gas discharge tube that
uses mercury as the gas and a phosphor coating on the inside of the tube. Mercury produces UV radiation that is at the same
energy that the phosphor absorbs and excites the phosphor electrons several energy levels up. The phosphor will reemit in the
visible wavelengths as the excited electrons in the phosphor drop in steps to lower energy levels. Not only do you get the
mercury visible emission lines but also the visible wavelengths from the phosphor which increases the intensity of visible
light. The light of the fluorescent tube is white because it has color bands of roughly equal intensity in all the color region
(but a little more on the blue side) where sunlight is more yellowish because it has a little more intensity in the red yellow
region. Since fluorescent bulbs produce an intense white light that uses less electricity (less of the energy goes into heat) it
has become quite popular. One downside of fluorescent bulbs is that cosmetics, when viewed beneath a fluorescent bulb, will
not appear the same as in sunlight which has more yellow in it. Another more long term downside is that the mercury in the
bulbs, when disposed of in landfills, can seep into groundwater and is toxic to plants, animals and humans. Mercury, like
lead, tends to accumulate in the body tissues over an organisms lifetime.
Historical Connections
The work of Max Planck, Albert Einstein and Niels Bohr lead the way to explaining that electrons exist in atoms in fixed
energy orbits called the Bohr model (simple, only works with one electron systems). These orbits were said to be “quantized”
because the energy differences were fixed and the electrons could only exist at these energies and not anywhere between.
Later after Louis DeBroglie and Heisenberg proposed that moving electrons behaving like waves, with the orbit path
described by probability, the quantum mechanical model was developed (sophisticated, could work with any number of
electrons but mathematically more difficult with increasing numbers of electrons). In quantum mechanical model the orbits
became orbitals at fixed energies and the electron path is described as a cloud of electron density that is in interesting shapes
like spheres (s), dumbbells (p), and clover leafs (d) depending on the level it was on. Pauli later determined that only two
electrons could fit in any orbital. When an element’s electrons are at the lowest energy that they can be (in the lowest orbitals)
they are said to be in the ground state (kind of like resting). In both the Bohr model and the quantum mechanical model
electrons could only jump up to higher energies if they had the exact energy supplied that equaled the energy difference
between orbitals. This could be done in two ways: thermally heating the element or by absorbing a photon of exactly that
energy. When elements have electrons that have jumped up to higher energy orbitals the electrons are said to be in the excited
state. These electrons will eventually fall back down to lower orbitals until they are at the lowest available orbitals that aren’t
General Chemistry
Lab 6: Spectroscopy
3
already filled,. Each time that an electron falls to a lower orbital it emits a photon that is exactly the energy difference of the
two orbitals. In this way we have conservation of energy.
n=4
n=3
a
Energy
Figure 3 shows two possible paths for an electron that has been
excited to level 3 in the Bohr model. First is path “a” where the
electron has dropped from level 3 to level 1 directly (solid arrow) and
emits a photon (dashed arrow) of that same energy difference. This
will be abbreviated as n=3Æn=1. The other path b + c will emit two
photons as the electron in level 3 first falls to level 2 shown as
n=3Æn=2 and emitting a photon then it falls to level 1 shown as
n=2Æn=1 emitting the second photon. Since the diagram in Figure 1
represents energy it should be obvious that the energy of the photon
emitted from drop “a” is greater than from drop “b” which is greater
from drop “c”. But you can visually see that the energy drop of “a” is
equal to the energy drop of “b + c”. Notice that as the energy increases
the levels get closer together.
b
n=2
c
n=1
Figure 3: The transitions in the hydrogen atom as
proposed by Bohr. An electron that jumped to the
n= 3 level had two paths down: n=3Æn=1
OR by two steps, n=3Æn=2 then n=2Æn=1
In today’s lab, you will be determining the energy associated with the various colors that you observe. You will first make a
calibration curve using your measured data correlated to the known wavelengths from hydrogen gas. You will then use this
curve to determine the wavelengths of the other lines that you observe with the other gases. Once you have the wavelength,
you can calculate the energy of each line
Mathematical Formulas
The wavelength of light produced by an electron determines the light's color, and the color is an indicator of the energy of the
orbit transition. If the wavelength of light can be measured, then the energy of that wavelength may be calculated via the
following equation from the laws of physics:
Energy (E) is Planck's constant (h) times the velocity of light, (c)
over the wavelength, (λ)
Planck's constant has a known value………………………………………
The speed of light also is known…………………………………………….
Putting everything together:
Energy =
h•c
λ
h = 6.63 x10–34 Js (joules seconds)
c = 2.998 x108 m/s
(6.63x10 −34 J ⋅ s ) • (2.998 x108
λ (in units of meters)
E=
m
)
s
General Chemistry
Lab 6: Spectroscopy
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Procedure
My lab partner is: ________________________________________
Write down detailed qualitative and quantitative observations in ink directly on the data collection sheet following this
procedure.
Equipment Setup
Spectroscopy is the study of how electromagnetic radiation interacts
with matter. A spectrometer is used to measure these interactions.
Light may contain radiation of many different wavelengths. The
appearance of a rainbow reveals that plain sunlight is actually a
mixture of many colors. Droplets of water in the atmosphere act as
thousands of prisms to break sunlight up into its component colors.
Spectroscopy labs use a diffraction grating to do the same thing.
In this experiment, light from an electric lamp (A) is viewed through
a diffraction grating, (B) . The grating separates the light so the
component colors may then be viewed against a meter stick (C).
Figure 4: Light from the light box is diffracted by the diffraction
grating into the component colors.
Figure 5: Lab setup of lamp (A), diffraction
grating (B), and meter stick (C).
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Lab 6: Spectroscopy
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Data Collection
1.
2.
3.
4.
5.
Obtain a light box, diffraction grating, grating holder, 2 meter sticks with two stands and pointer
and set them up according to the diagram in Fig 2 on the first page.
Put the meter stick (B) on stands and tape it down so the enc of the meter stick is at the end of the bench. Put the grating
holder near the end of the meter stick and tape it down. Put the diffraction grating on the holder (make sure that color
images go horizontal not vertical). (see fig 2)
Make the light box be as far back as possible. Mark the position of the light box with chalk or tape to keep it in one place
(it should not move when you change the bulbs).
The meter stick (C) should be placed flat and in front of the light box (for illumination in dark). You may want to let it
overhang the table a little. Make sure the meter stick is taped into place to prevent it from moving
For each of the tubes: hydrogen, helium, mercury, neon
a. Select a light tube and install it in the light box. Handle the glass tubes carefully.
b. Use caution when changing the tubes: they can get hot. Use a dry paper towel to protect your fingers.
c. Turn on the light box and look through the diffraction grating. You should see several brightly colored lines.
d. One partner should be in place to mark the virtual position of the lines, as directed by the person viewing. Make sure
your partner also has a chance to observe spectra for each element.
e. Describe and record on the data sheet on page 6: the color of each distinct line and the position of its image on the
meter stick. Omit blurry images and ignore the overtones that occur further away from the first set of lines.
Note: You are viewing line spectra that should have distinct colored lines and a BLACK background. Because of dirty
tubes (oils from hands), old tubes and dirty diffraction gratings you may see a light rainbow smear in the background.
Only the bright lines should be there so ignore any colored background you see.
•
For hydrogen, obtain data for the three of the four standard lines. (The violet line is hard to see.)
•
For helium and mercury, choose up to six lines evenly distributed between purple and red.
•
6.
7.
For neon, group the lines according to color and estimate the number of lines in each group. Assign the
meter stick reading to be in the middle of the color group
View light from the fluorescent overhead lights and an incandescent source (light bulb) from about 10 feet away. Record
your description to compare a continuous spectrum with the line spectra observed from the gas tubes. Note any strong
(bright) lines in the spectrum that may be present.
On the bottom of the page put the smallest and largest meter stick values in cm for all the data taken.
After you have collected all the data keep your set up together just in case you need to re-measure.
General Chemistry
Table 1: Direct Data Collection
Element
Line Color
Hydrogen
violet (434nm)
teal (486nm)
red (656nm)
Lab 6: Spectroscopy
Date ___________________________
Position (cm)
on the meter stick
Description or Notes
Helium
Mercury
Neon
Description:
Fluorescent
light source
Description:
Incandescent
light source
For ALL the measurements above: smallest distance (cm) _______________ largest distance (cm) _______________
(These values are used to make the graph y-scale: your data extremes should fill greater than 1/2 the graph)
Remember:
• For hydrogen, obtain data for the three (3) standard lines. The violet line is hard to see.
• For helium and mercury, choose four (4) lines evenly distributed between purple and red.
• For neon, group the lines according to color and estimate the number of lines in each group. Assign the meter stick
reading to be in the middle of the color group.
6
General Chemistry
Lab 6: Spectroscopy
7
Data Analysis
For calculations, SHOW ALL WORK NEATLY FOR ONE SAMPLE CALCULATION. Use proper units and unit
conversions throughout each calculation. Report your answer with correct significant figures.
Determining wavelength and energy for the line spectra of elements observed
A. Make a Calibration Graph. This is a translator that takes your readings (cm) and changes them to wavelengths in
nanometers.
Prepare a graph of wavelength in nm (x-axis) versus diffraction distance in cm (y-axis) for HYDROGEN ONLY
since you are given the wavelengths of the hydrogen lines.
1. Using the highest and lowest “cm” values to create y-axis scale (have instructor help).
2. Plot the data points you obtained from the Hydrogen lamp. The known colors and wavelengths for hydrogen are
as follows. Violet: 434nm Teal: 486 nm
Red: 656 nm
3. Draw a "best fit" straight line through the points. DO NOT connect the dots.
This is known as a calibration curve. In part B you will use the graph to convert distance measured in cm on the
meter stick to wavelength in nanometers for all lines you recorded.
Graph of Line Position vs. Wavelength (Based on Hydrogen data)
(use the largest and smallest cm values from all the spectra to create the vertical y-scale)
Position of lines on meter stick (cm)
I.
400
450
500
550
Wavelength, λ. (nm)
600
650
700
B. Determining Wavelength from Calibration Graph and Calculation of Energy
Now that we have a standard for comparison (the calibration curve), we can use it to determine wavelength from the
measured position on the meter stick.
1. For every line measured, mark and label the diffraction distance point (cm measurement from meter stick) on
the y-axis.
2. Use a ruler to trace a horizontal line from the mark to the calibration curve.
3. Where the line intersects the calibration curve, use a ruler to trace a line down to the x-axis.
4. Read the wavelength in nm where the line intersects the x-axis, and record in Table 2.
5. From the wavelength, calculate the energy of each line and record this in Table 2
[Check: for Hydrogen: Violet ( 434 nm) the energy should be E = 4.58x10-19 J]
Suggestion: The numerator of the energy calculation ( h x c ) is the same for all lines. Enter this number into
memory in your calculator; then recall the number and divide by lambda for each line.
6. Show one sample calculation for determining energy (below Table 2 on the next page).
Check with instructor on significant digits and units if you have questions.
General Chemistry
Lab 6: Spectroscopy
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TABLE 2: Wavelength and Energy Determination
(First copy color and cm position from Table 1)
Element
1.
Line Color
Show sample calculation: Energy =
cm position on
meter stick
Wavelength
( nm )
Energy in Joules
(calculation, E=hc/λ)
ranking
General Chemistry
Lab 6: Spectroscopy
II. Relationship of Energy to Rainbow Colors
1.
2.
On the side of Table 2 rank the energies (1, 2, 3, …) from lowest to highest where 1 is lowest.
In Table 3 below, arrange ALL the line spectral data in Table 2 by energy with the lowest energy on top (use your
rankings to help).
TABLE 3. Emission Lines Ranked by Energy .
Arrange the lines from Table 2 according to increasing energy, 1=lowest energy.
Rank
Element
Color
Wavelength( nm)
Energy (Joules)
9
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Lab 6: Spectroscopy
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III. Relationship of Observed data to Reference Data
You probably noticed that some lines you observed today are very intense and some are very faint . The following is a list of
reference wavelengths for various elements, including those seen in our lab today. Some of the lines are very faint and can
only be seen in a darkened room with very good eyes. Sometimes a gas tube gets leaks, is dirty, or gets old and some new
lines are present.
TABLE 4 Reference Line Spectra at Visible Wavelengths (nm)
Calcium
393
397
423
Helium
447
469
502
588
688
Hydrogen
410
434
486
656
Iron
431
438
467
496
527
Oxygen
686-688
Mercury
405
436
546
579
615
691
Sodium
589
590
Neon
470
535
585
640
694
718
a.
In the above table, for each element you saw today, circle wavelengths that match the closest to every wavelength that
you measured today (3 for H, 4 for He, etc.) even if 20 nm away.
b.
Emission line spectra are unique for each element (like a fingerprint). From the values in the table above only, do any
values from different elements match exactly? Which sets of emission lines are closest? Explain.
c.
If you burned a compound and noticed that the wavelengths 438, 467 and 527 were present in the flame what element
might you suspect is present (circle the element name)?
Put all items back on the cart. Return the key to the front desk. Obtain instructor initials before leaving the lab.
Questions
Instructor Initials ___________
Answer the questions below on a separate piece of paper. If you answer questions on this page your lab report will be
penalized 5%. We want to know if you can follow directions.
Answer on separate paper and attach to report. State the question in the answer, or write out the question.
1. In the data collection, were all four lines for hydrogen easy to see? Elaborate.
2. What does ROY G. BIV represent? Does the order in your table (Table 3) follow the order of colors in the rainbow?
(ROY G. BIV). Explain why any colors are out of order.
3. Did your experimental wavelengths in general fall lower, higher, or evenly lower and higher than the reference values
listed in part III of the Data Analysis?
4. Comment on the reliability of your results, including sources of experimental uncertainty.
5. How could you determine whether calcium is in a distant star?
6. Relate the principles learned to some aspect of daily life.
Summary Writing Assignment
Write the following in grammatically correct English sentences. Answer on a separate sheet of paper.
A. Answer the Experimental Question: How can the wavelength and energy of emission lines from various elements be
determined from the known wavelengths of a hydrogen spectrum? Explain using evidence from your experimental data.
B. Describe any points about today’s lab that still need clarification for you.
Turn in this Entire Packet PLUS attached sheets for
Questions and Summary Writing Assignment stapled to the back.