Chemistry lab 3

CHEMISTRY
Atomic Emission and
Flame Test
Investigation
Manual
ATOMIC EMISSION AND FLAME TEST
Table of Contents
2
Overview
2
Outcomes
2
Time Requirements
3
Background
6
Materials
7
Safety
7
Preparation
8
Activity 1
8
Activity 2
8
Activity 3
9
Activity 4
9
Disposal and Cleanup
Overview
In this investigation students will observe the color of light emitted
from excited metal ions. This information will be used to partially
determine the identity of an unknown salt. Each salt provided
has a unique color to help with metal determination. Students will
observe household light sources and determine the spectrum of
the light emitted. The unknown metal spectrum will be identified
by comparing it with the known metal spectrums.
Outcomes
• Conduct flame tests of metal ions.
• Determine the light wavelengths associated with an observed
spectrum.
• Determine the partial identity of an unknown salt using flame test
light emission.
• Create a calibration curve.
• Use a spectroscope to measure light emission.
• Calculate initial quantum number of hydrogen based on Bohr’s
equation.
Time Requirements
Preparation ………………………………………………………………5 minutes
Activity 1: Household Light Sources ………………………….10 minutes
Activity 2: Spectroscope Calibration ………………………….15 minutes
Activity 3: Metal Light Emission ………………………………..50 minutes
Activity 4: Calculation of Hydrogen Quantum Numbers .20 minutes
Cleanup …………………………………………………………………..5 minutes
Key
Personal protective
equipment
(PPE)
goggles gloves apron
Made ADA compliant by
NetCentric Technologies using
the CommonLook® software
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follow
link to
video
photograph stopwatch
results and
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submit
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Background
Many holidays are celebrated with fireworks
that illuminate the sky with a variety of colored
lights. The metal ions in fireworks are chosen
specifically to create the assortment of observed
colors. This colored light is produced by excitation and subsequent relaxation of the electrons
in the metals.
Light emission by atoms is a foundational principle of quantum chemistry. Atoms contain
electrons that occupy discrete energy levels. The
actual energy of each state (level) depends upon
several factors: the nuclear charge, the distance
of the electron from the nucleus, and the number
of other electrons between the nucleus and the
electron of interest.
Bohr’s Model of the Hydrogen Atom
Niels Bohr made one of the first large breakthroughs in atom modeling and light quantization using the spectrum of hydrogen. Bohr
developed a theory that electrons could only
occupy certain distinct energy levels around the
hydrogen nucleus. An example of these quantized energy states are shown in Figure 1.
Figure 1.
To transition an electron from one level to another, a discrete amount of energy must be emitted
or absorbed. Hydrogen is the easiest to observe
because it only contains one electron. For light
emitted by hydrogen, the energy can be determined as follows.
ℎc
ΔE =
λ
Where λ is the wavelength of the light in meters,
c is the speed of light (3.00 × 108 meters per
second), h is Planck’s constant (6.63 × 10-34
joules seconds), and ΔE is the change in energy
in joules.
This energy can be used to identify the quantum
number of the energy level the electron occupied. Quantum numbers identify the discrete
energy levels of electrons. Quantum numbers
are always integers. In hydrogen they start at
1, which is considered the ground state. If an
electron is in an orbital with a quantum number higher than 1, it is considered to be in the
excited state. There is more than one excited
state and electrons can relax between individual
excited states. In fact, with hydrogen visible light
is only emitted when the electron relaxes to an
orbital with a quantum number of 2. For visible
light, the electron’s quantum number prior to
relaxing to the n = 2 state can be calculated as
follows.
1
1
ΔE = (2.18 × 10−18J) ( n2
n2i )
f
Where ΔE is the change in energy calculated
earlier, nf is the quantum number of the final
orbital, and ni is the quantum number of the
initial orbital. Since hydrogen’s visible spectral
continued on next page
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ATOMIC EMISSION AND FLAME TEST
Background continued
lines all have an nf = 2, the initial quantum number is easily calculated. These calculations will
be performed in Activity 4. Bohr’s calculations
for quantum numbers only work for hydrogen.
For other elements electron configurations
become increasingly complicated and Bohr’s
equations are insufficient. Excited electrons
still emit photons and the energy’s magnitude
depends on the energy of the levels being transitioned between. The number and type of transitions depend both on the energy level structure
of a given chemical species and on various
quantum selection rules. These properties are
unique to each element and give rise to characteristic element emission energies. Since each
electron can undergo many transitions and most
elements have many electrons, emission spectra
can consist of a very large number of discrete
wavelengths.
These wavelength profiles can be leveraged
for scientific pursuits and practical effects. For
example, sodium emits light at 589.3 nm. Areas
surrounding astronomical observatories use low
pressure sodium street lights to minimize light
pollution. The observatories can filter out light at
specific wavelengths and still produce credible,
reproducible results even in dense urban areas.
At your next opportunity, look at some nearby
street lights at night. If they emit a strong yellow
light, there is a good chance they are sodium
lights.
The majority of light we observe is from
continuous emission spectra. This light contains
so many emitted radiation frequencies that the
lines overlap. The light appears white, but it is
really a rainbow of colors. Common continuous
emission sources are incandescent light bulbs,
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heated metals in general, and flames containing soot such as a butane lighter. Notice that all
these sources are solids, e.g., the wire filament
in a light bulb and soot particles in a candle
flame. The atoms and molecules in heated solids are continuously bumping against each other
because they are so close. This dynamic contact
causes the excitation energy to transform into
kinetic energy, which is not quantized. As a solid’s temperature is raised, more of the radiation
is emitted at shorter wavelengths. One example
of this phenomenon is an electric stove. When
the stove is first turned on, the coils warm and
emit infrared radiation, which is emission at
wavelengths greater than 700 nm. As the coils
continue to heat up, red light is emitted. Most
stoves do not heat above the red hot stage, but
if it continued the coil would appear orange,
then yellow, and eventually white as the
temperature increased.
Using a spectroscope
A spectroscope is an instrument that measures
the wavelengths of light emitted from a light
source. The spectroscope provided in this kit
has three main components. There is a slit at
one end of the spectroscope that allows light
to enter the instrument. The light then travels to
the diffraction grating where the light is reflected
and refracted towards the scale at the far end
of the spectroscope. The diffraction grating acts
like hundreds of prisms, bending the light such
that the colors are split and focused on the scale
in discrete wavelengths. The numbers on the
scale are roughly equivalent to the associated
light wavelength, e.g., a 4 on the scale is equivalent to a 400 nm light wavelength.
continued on next page
Figure 2.
To use the spectroscope, look through the
diffraction grating and align the slit at the far end
of the spectroscope with the light source being
observed. Once the light source is properly
aligned, the light spectrum will appear on the
scale in the spectroscope. This light may be very
faint and difficult to observe even under ideal
circumstances.
Table 1.
Color Wavelength
(nm)
Violet
400 – 430
Blue
430 – 490
Green
490 – 570
Yellow
570 – 590
Orange
590 – 640
Red
640 – 750
In Activity 2, you will use a fluorescent light
spectrum to calibrate your spectroscope.
Modern fluorescent bulbs use a combination
of elements to produce visible light. The most
common elements are listed in the table below.
Table 2.
Peak
Number
Peak
Wavelength
(nm)
Peak
Color
Element
1
405.4
Purple
Mercury
2
436.6
Blue
Mercury
3
487.7
Blue green
/Teal
Terbium
4
546.5
Green
Mercury
5
625.7
Orange
Terbium
6
662.6
Red
Europium
The calibration in Activity 2 is then used to
directly correlate the scale position observed in
the spectroscope with a particular light wavelength. Activity 3 will use the spectroscope and
the calibration curve constructed in Activity 2
to determine the prominent spectral lines and
approximate wavelengths for four metal ions
and an unknown metal ion mixture that you will
identify.
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ATOMIC EMISSION AND FLAME TEST
Materials
Included in the materials kit:
Needed from the equipment kit:
Sterno® can
Plastic cup,
10 oz
Flame Test Sticks:
KCl – Purple color
LiCl – Red color
CuCl2 – Blue-green color
NaCl – Orange color
Unknown salt – colorless
Spectroscope
Reorder Information: Replacement
investigation kit for Atomic Emission and
Flame Test, item number 580350, can be
ordered from Carolina Biological Supply
Company.
Call 800-334-5551 to order.
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Forceps
Needed but not supplied:
• Matches or lighter
• Digital camera or smart phone
• Graphing program (e.g., Excel®)
• Oven mitt or tongs
• Trivet or cooling rack (optional)
Safety
Read all the instructions for this laboratory
activity before beginning. Follow the
instructions closely and observe established
laboratory safety practices including
the use of appropriate personal protective
equipment (PPE) as described in the Safety
and Procedure sections.
Preparation
1. Obtain all materials.
2. Select a work area with a heat-resistant
surface, preferably a stovetop or trivet.
3. Remove any flammable household materials
from the work area.
Wear your safety goggles
and gloves at all times
while conducting this
investigation.
Sterno® contains flammable alcohol
gel. Long hair should be tied back
and loose clothing kept away from an
open flame. The can will stay hot even
after the flame is extinguished. Use
Sterno® on a heat-resistant surface like
a stovetop or trivet.
The potassium, lithium, sodium, and
copper metal salts used are toxic when
ingested. Handle with gloves at all
time. Keep flame test sticks away from
food storage locations, children, and
pets to avoid accidental ingestion.
Do not eat, drink, or chew gum while performing
this activity. Wash your hands with soap and
water before and after performing the activity.
Clean up the work area with soap and water
after completing the investigation. Keep pets
and children away from lab materials and
equipment.
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ACTIVITY
ACTIVITY 1
ACTIVITY 3
A Household Light Sources
A Metal Light Emission
1. Using the spectroscope, view a fluorescent
light bulb.
1. This activity works best in a darkened space.
If possible, have an assistant to hold the
coated sticks in the flame, so that observing
the color of the flame is easier.
2. Record the color and scale position of the
three strongest spectrum lines in Data Table 1.
Start on the left and proceed to the right.
3. Observe two additional light sources and
record whether the light has a continuous or
line spectrum in Data Table 1.
4. If the light has a line spectrum record the
color and scale position of the spectrum lines
in Data Table 1.
Different types of lights will likely have
different spectra. Feel free to be creative and
look at different types of light sources in the
world around you.
ACTIVITY 2
A Spectroscope Calibration
1. Copy the scale position and color for each
observed line from the fluorescent light bulb
into Data Table 2. The number of lines may
vary based on the particular light source.
2. Reference the Common Elements table in
the Background section and write the line’s
wavelength in Data Table 2.
3. Use a graphing program to create a scatterplot graph with the light wavelength on the
y-axis and the scale position on the x-axis.
4. On the graph, create a best-fit line (linear
trendline) for your scatter plots.
5. Set the y-intercept of the best-fit line to 0.
6. Display the trendline equation on your graph
and record it in Data Table 2.
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2. Fill a plastic cup with tap water.
3. Open the Sterno® can by gently prying up the
lid with a key or knife.
4. Light the Sterno® heater with a match or
lighter. A Sterno® flame is nearly invisible in a
bright room and appears pale blue when the
lights are darkened.
5. Turn off any interior lights and make the
working area as dark as you can while still
feeling comfortable performing the task.
The Sterno® heater will get extremely hot
as the activity progresses. It is important
to select a suitable heat-resistant working
area so you do not need to move the can.
Keep the lid and oven mitts or tongs handy.
The Sterno® lid can be slid over the flame to
immediately extinguish the can. Sterno® can
be relit if extinguished.
6. Take the flame test stick coated in potassium
chloride (KCl) and break it into four pieces.
Refer to the materials section for help
identifying the test sticks by color.
7. With the forceps, carefully dip one piece
of the flame test stick into the water in the
plastic cup. This prevents the wood burning
and influencing the color of the flame.
8. With the forceps, carefully hold the flame test
stick over the Sterno® heater. The wood will
burn quickly then reveal the unique colored
flame of the metal ion.
ACTIVITY 3 continued
9. Record the color and intensity of the stick’s
flame in Data Table 3.
Disposal and Cleanup
1. With an oven mitt or tongs, slide the lid over
the Sterno® flame to extinguish.
10. If necessary, add additional pieces of flame
test stick to the can for prolonged or brighter
color.
2. Once cool, the Sterno® can and remaining
flame test sticks can be disposed in a waste
receptacle.
11. View the flame through your spectroscope
and record any spectrum line(s) observed in
Data Table 3.
3. Sanitize the workspace.
12. Repeat steps 5 through 11 with the
remaining salts in order (LiCl, NaCl, CuCl2,
and the unknown salt).
13. Calculate the light wavelength for each line
observed in all the compounds. Use the best
fit line equation determined in Activity 2 to
perform this calculation. Replace x with the
observed scale position and solve for y.
ACTIVITY 4
A Calculation of Hydrogen Quantum
Numbers
1. In Data Table 4, four wavelengths of light were
observed in the hydrogen light spectrum.
2. Convert the wavelengths from nanometers to
meters and record the values in Data Table 4.
3. Calculate ΔE for each wavelength and record
the values in Data Table 4.
4. Calculate the initial quantum number, ni, for
each spectral line and record the values in
Data Table 4.
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DATA TABLES
Data Table 1: Household Light Sources
Light Source 1
Name
Fluorescent Bulb
Line or Continuous?
First Spectrum Line
Color
First Spectrum Line
Number
Second Spectrum
Line Color
Second Spectrum
Line Number
Third Spectrum Line
Color
Third Spectrum Line
Number
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Light Source 2
Light Source 3
Data Table 2: Spectroscope Calibration
Line Color
Line Scale Number
Wavelength (based
on Table 2 in the lab
background)
First Spectrum Line
Second
Spectrum Line
Third
Spectrum Line
Fourth
Spectrum Line (if visible)
Fifth
Spectrum Line (if visible)
Sixth
Spectrum Line (if visible)
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DATA TABLES
Data Table 3: Metal Salt Light Emission
Metal
Salt
Potassium
Chloride
Color
of Flame
First Spectrum
Line Color
First Spectrum
Line Number
First Spectrum
Line Wavelength
Second Spectrum
Line Color
Second Spectrum
Line Number
Second Spectrum
Line Wavelength
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Lithium
Chloride
Copper
Chloride
Sodium
Chloride
Unknown
Salt
Data Table 4: Calculation of Hydrogen Quantum Numbers
Constants needed for the calculations:
c = 3.00 x 108 meters/second
Planck’s Constant (h) = 6.63 x 10-34 joules seconds
nf = 2
Wavelength
(nm)
Red
Light
Green
Light
Blue
Light
Violet
Light
656.2 nm
486.1 nm
434.0 nm
410.0 nm
Wavelength
(m)
ΔE (J)
ni
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NOTES
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NOTES
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CHEMISTRY
Atomic Emission and Flame Test
Investigation Manual
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866.332.4478
Carolina Biological Supply Company
www.carolina.com • 800.334.5551
©2017 Carolina Biological Supply Company
CB780861703
Atomic Emission and Flame Test
Student Name
Date
1
Data
Activity 1
Data Table 1
Name
Light Source 1:
Fluorescent Bulb
Line or
Continuous?
First Spectrum
Line Color
First Spectrum
Line Number
Second Spectrum
Line Color
Second Spectrum
Line Number
Third Spectrum
Line Color
Third Spectrum
Line Number
© 2016 Carolina Biological Supply Company
Light Source 2:
Light Source 3:
2
Activity 2
Data Table 2
Line Color
Line Scale
Number
Wavelength (based on Table
2 in the lab background)
First Spectrum
Line
Second
Spectrum Line
Third Spectrum
Line
Fourth Spectrum
Line (if visible)
Fifth Spectrum
Line (if visible)
Sixth Spectrum
Line (if visible)
Graph
Insert your graph here:
1. What is the equation for the linear trend-line for your equation?
2. What is the corrected equation for the trend-line?
© 2016 Carolina Biological Supply Company
3
Activity 3
Data Table 3
Metal Salt
Potassium
Chloride
Lithium
Chloride
Copper
Chloride
Sodium
Chloride
Unknown
Salt
Color of
Flame
First
Spectrum
Line Color
First
Spectrum
Line Number
First
Spectrum
Line
Wavelength
Second
Spectrum
Line Color
Second
Spectrum
Line Number
Second
Spectrum
Line
Wavelength
3. What wavelengths are associated with the light color observed in each
metal’s flame test?
4. What are the most probable atoms in your unknown solid based on the
observed light emissions?
5. Sodium light is easy to filter because it only emits light at a specific
wavelength. Would any of the other metals you tested also be easy to filter?
Why or why not?
6. If a fireworks engineer was planning a Christmas show and wanted to create
a red explosion followed by a green one. Based on the flames observed,
which metal salts should they use?
© 2016 Carolina Biological Supply Company
4
Activity 4
Data Table 4
Constants needed for the calculations:
c = 3.00 x 108 meters/second
Planck’s Constant (h) = 6.63 x 10-34 joules seconds
nf = 2
Wavelength
(nm)
Wavelength
(m)
Red Light
Green Light
Blue Light
Violet Light
656.2 nm
486.1 nm
434.0 nm
410.0 nm
ΔE (J)
ni
© 2016 Carolina Biological Supply Company