Optical Spectroscopy

Objective:
Observe emission spectra from a variety of sources using a spectrometer
Design and implement a procedure capable of determining the composition of a solution
that contains two or more ionic salts
Construct a partial energy-level diagram for hydrogen
Introduction:
Optical spectroscopy involves the measurement and analysis of electromagnetic radiation
(a.k.a. light). Many of the properties of light are conveniently described by means of a classical
wave model. Within this model, light waves are characterized by such variables as frequency and
wavelength. The frequency (), which describes the number of wave crests passing a given point
per second for the light wave, is inversely proportional to the wavelength (), the distance
between successive wave crests. The light we can see, visible light, corresponds to a very small
portion of the electromagnetic spectrum, from about 400 nm (violet) to 800 nm (red). The light
wave frequency and wavelength are related to one another by the equation
c =

(1)
8

where c is the speed of the light wave. The speed of light in a vacuum is 2.998 10 m/s.
For phenomena where the classical wave model of light proves insufficient, a particle
model is invoked. In the particle model, light is composed of a stream of discrete particles called
photons. The energy (E) of each photon is directly proportional to its frequency and inversely
proportional to its wavelength,
hc
E = h =
(2)

where the proportionality constant h is Plancks constant (6.62611034 Js). Therefore,

electromagnetic radiation can be described by either the wave or particle model; the model
applied is the one that accurately describes the phenomenon being investigated.
In spectroscopy, light is used as a means of probing matter. One means of probing matter
with light uses the phenomenon of absorption. When an atom, molecule, or ion absorbs a photon,
its energy increases. The energy change of the atom must be equivalent to the energy of the
photon. Thus the absorbed wavelengths of light reveal the differences between energy levels:

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Optical Spectroscopy
E3 = 8.0 aJ

E3,2

E2 = 17.0 aJ

E3,1 = E3 E1 = 31.5 aJ
1 = hc/E3,1 = 6.31 nm
E2,1 = E2 E1 = 22.5 aJ
2 = hc/E2,1 = 8.83 nm

E3,1
E2,1

E3,2 = E3 E2 = 9.0 aJ
3 = hc/E3,2 = 22 nm

E1 = 39.5 aJ
1

Figure 1: For an atom with only three energy states as shown above, there are
only three wavelengths of light that can be absorbed, each illustrated by an arrow.
The values of the three absorbed wavelengths of light, as calculated from the
energy-level differences, are shown as well.
The converse of absorption is emission; when an atom emits a photon of light, its energy
decreases. The energy of the emitted photon must be equivalent to the energy change of the
atom. The emitted wavelengths of light correspond to the differences between energy levels:
E3 = 1.5 aJ

E2,3

E2 = 16.0 aJ

E1,3 = E1 E3 = 43.5 aJ
1 = hc/E1,3 = 4.56 nm
E1,2 = E1 E2 = 29.0 aJ
2 = hc/E1,2 = 6.85 nm

E1,3
E1,2

E2,3 = E2 E3 = 14.5 aJ
3 = hc/E2,3 = 13.7 nm

E1 = 45.0 aJ
1

Figure 2: Analogous with the description in Figure 1, only three wavelengths of

light can be emitted, each illustrated by an arrow. Negative energies arise from
the thermodynamic convention that energy released by an atom is represented by
negative E values; the values of the emitted wavelengths calculated from
energy-level differences are shown to the right of the energy-level diagram.
In many cases, the pattern of wavelengths absorbed or emitted by a pure substance is
characteristic of that substance. Thus the pattern of emitted or absorbed wavelengths can be used
as a means of identifying a substance, a sort of fingerprint in light.
The patterns of wavelengths absorbed or emitted by atoms, molecules, or ions are known
as spectra. Spectra may be classified as emission spectra or absorption spectra. In an emission
experiment, the source, which could be an ordinary incandescent or fluorescent light bulb, a salt
in a flame, or an electrically excited gas in a tube, emits the light (see Figure 3).

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Optical Spectroscopy

Source

Many s

Wavelength
selector

Detector

One

Figure 3: Typical emission experimental set-up.

The emitted light is passed through a wavelength selector (a prism or diffraction grating) to
select one wavelength. The intensity of light at this wavelength is measured at the detector.
Adjusting the wavelength selector changes the wavelength of light whose intensity is measured at
the detector. The collection of these measurements over a range of wavelengths makes up the
spectrum. If only a few characteristic wavelengths are emitted, the result is a bright-line emission
spectrum. If emission occurs at all wavelengths within a given range, the result is a continuous
emission spectrum. Continuous emission spectra result when the number of available energy
levels is very large and the spacing between them approaches the infinitesimal.
In an absorption experiment, the light from a source passes through an absorbing
medium, such as a gas sample or a solution, the effect of which is to remove certain wavelengths.
The wavelengths of unabsorbed light then are passed through a wavelength selector and onto a
detector (see Figure 4).
Source

Many s

Absorbing
medium

Fewer s

Wavelength
selector

One

Detector

Figure 4: Typical absorption experimental set-up.

The result is a dark-line spectrum. If a band of wavelengths is absorbed, the result will be
a dark area in that part of the spectrum.
In this experiment you will be using both a simple hand-held spectroscope and a
sophisticated research-grade spectrometer to acquire emission spectra from a variety of sources.
Both of these instruments contain a diffraction grating (wavelength selector) and a detector. The
hand-held spectroscope presents this data in a manner that is more visually appealing, as colors
on a wavelength scale. The research-grade spectrometer is more quantitative; it yields
quantitative measures of intensity as a function of wavelength. The different views of emission
spectra obtained by the two instruments complement one another.
The emission spectra of a variety of materials will be observed: a fluorescent light bulb, an
incandescent light bulb, helium gas, and a variety of salt solutions. These materials will provide
examples of both continuous and line spectra.
Your purpose in this experiment is twofold. One goal is to determine the composition of a
solution that contains two or more ionic salts. The second goal is to construct a partial energylevel diagram for hydrogen (see Figures 1 and 2 for examples of this). In Part A of this
experiment, you will make a number of observations of emission spectra. In Part B, you will use
your initial observations to design procedures for determining the composition of a solution that
contains two or more ionic salts and for constructing a partial energy-level diagram for
hydrogen.

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Safety Precautions:
Safety goggles must be worn at all times while you are in the laboratory.
In addition to visible light, the discharge tubes emit ultraviolet radiation, which is
damaging to the eyes. While safety goggles will absorb most of this radiation, it is
Experiment:
recommended that you look at the radiation source for only short periods of time.
When
recording
data, include
thedevelops
relative intensity
line.
If your volts.
The
power
supplyspectral
to the discharge
tubes
a voltageof
ofeach
several
thousand
spectrometer
intensity
the intensities
by eye
a scale
of 1unless
to 10,
Do notlacks
touchanany
portionscale,
of theestimate
power supply,
wire leads,
or using
discharge
tubes
where 10
means
very
bright
and
1
means
you
can
barely
see
the
line.
For
continuous
spectra,
the power supply is unplugged from the electrical outlet.
record
ranges
to each
that you
canprior
distinguish.
wavelength
Always unplug
thecorresponding
power supply from
thecolor
electrical
outlet
to adjusting the
of the discharge
tubes
or any other part of the apparatus.
Part A:position
Observations
of emission
spectra.
Procedure:
Part A: Observation of Emission Spectra
1. Fluorescent Light Spectrum:
Observe the emission spectrum of the fluorescent light with the spectroscope and the
spectrometer. Some of the lines you are likely to see occur at 405, 436, 546, 577, 579, 615, and
691 nm. Dont be concerned if you cant see all the lines; some are very faint.
2. Incandescent Light bulb Spectrum:
Observe the emission spectrum of the incandescent light bulb with the spectroscope and
the spectrometer. You should observe a continuous spectrum because an incandescent light bulb
gives off white light.
3. Helium Spectrum:
WARNING: Because of the ultraviolet radiation emitted, you should look at the radiation source
for only short periods of time. Do not touch any portion of the power supply, wire leads, or
discharge tube unless the power supply is unplugged from the electrical outlet because of the
very large voltages produced by the apparatus.
There is a device in your lab consisting of a glass tube (discharge tube) filled with helium
that is attached to a voltage source used to excite the helium gas, causing it to glow. Turn the
discharge tube on, and record the color of the glowing helium gas. Use the spectroscope and the
spectrometer to observe the helium emission spectrum. Helium lines reportedly occur at 447,
502, 588, 668, and 707 nm, but some are much easier to see than others.
4. Spectra of Salt Solutions: NaCl, LiCl, KCl, CaCl2, and SrCl2
Half-fill five of the smallest test tubes in your equipment drawer with one each of the
following salt solutions: sodium chloride, lithium chloride, potassium chloride, calcium chloride,
and strontium chloride. Obtain 10 cotton swabs.
Light a Bunsen burner. Adjust the air flow using the knurled knob at the bottom of the
burner so as to get two distinct cones of flame. Arrange the spectrometer so that the flame is
visible in the slit.
Caution: Dont get the spectrometer too close to the flame; the heat could damage it.

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Use the spectrometer to observe the emission spectrum of each salt solution. To
accomplish this, soak one end of a cotton swab in a salt solution. One student should hold the tip
of the swab in the hottest portion of the burner flame (just above the inner cone) while another
looks through the spectroscope at the flame. It may take a few seconds for the swab to dry out
before the intensely colored flame appears. Try not to ignite the swab. If you cannot find the slit
image when the color is intense, soak the swab again, and repeat the trial. Repeat this procedure
for each of the other salts. To avoid contamination, be sure to use a new swab when you change
salt solutions.
The emission from potassium chloride is so faint that you may have difficulty seeing any
lines at all. If you cannot record the wavelength of any potassium lines, just record the color.
Part B: Analysis of Spectra
1. Identification of Unknown Salts in Solution:
You are to design and carry out a procedure that will allow for the identification of the
unknown salts in a solution. Your unknown solution will contain two or more of the following
salts: NaCl, LiCl, KCl, CaCl2, and SrCl2. Your instructor will provide you with your unknown.
Record the unknown number in your laboratory notebook.
Prior to examining its unknown sample, you are required to test your procedure on a
mixture of known composition. Use the salt solutions available in labNaCl, LiCl, KCl, CaCl2,
and SrCl2to create a mixture of known composition. Record the results of testing your
procedure. Use these data to verify (or improve) the efficacy of your procedure.
2. Partial Energy-Level Diagram for Hydrogen:
WARNING: Because of the ultraviolet radiation emitted, you should look at the radiation source
for only short periods of time. Do not touch any portion of the power supply, wire leads, or
discharge tube unless the power supply is unplugged from the electrical outlet because of the
very large voltages produced by the apparatus.
There is also a discharge tube filled with hydrogen in the lab. You are to design and
implement a procedure for collecting data from this discharge tube that can be used to generate a
partial energy-level diagram for the electronic states of hydrogen.
For the partial energy-level diagram of hydrogen, assume that all the observed transitions
terminate at the n = 2 state; for example, if you observe two transitions, they are from state A
n = 2 and state B n = 2. Also, set the value of the energy of the n = 2 state to 0.545 aJ.
Available Equipment and Reagents:
To perform this experiment, you will have access to all the equipment in your lab drawer
and:

cotton swabs
a spectrometer (equipped with a fiber optic cable) and a spectroscope
hydrogen and helium discharge tubes, incandescent and fluorescent lights
aqueous solutions of NaCl, LiCl, KCl, CaCl2, and SrCl2

Waste Disposal:
All chemical waste is to be flushed down the sink with plenty of water.
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Name: _________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
EXPERIMENT 7 Optical Spectroscopy
Pre-laboratory Questions: (answers to be written in your laboratory notebook)
1. Explain the difference between continuous and line spectra.
2. Explain the difference between absorption and emission spectra.
3. For an atom with the energy levels below, what wavelength light (in nm) will be emitted
in a transition between E5 and E1 (indicated by the down arrow below)? What wavelength
of light must be absorbed to cause a transition between E2 and E5 (indicated by the up
arrow below)?
E5 = 0.0133 aJ
E4 = 0.0282 aJ
E3 = 0.0656 aJ

E2 = 0.332 aJ

E1 = 1.827 aJ

4. Can the atom of Question 3 absorb or emit light with a wavelength of 723 nm? Can it
absorb or emit light with a wavelength of 653 nm? If so, state which energy levels the
transition occurs between. Show all work.
5. Consider the emission spectra of the two hypothetical elements X and Z. (NOTE: On the
spectra below, the upper scale corresponds to energies in units of eV.)

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Cut out these three spectra below and paste them in your laboratory notebook to make
answering this question easier. Since the duplicate copy will be collected, make sure that
the spectra appear on the duplicate copy.
Emission spectrum of X:
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Emission spectrum of Z:
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Draw a picture of the emission spectra expected from a sample containing a mixture of X
and Z on the spectrum blank below.
Emission spectrum of a mixture of X and Z:
1.7

1.8

1.9

700

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2.0

2.2

600

2.4

500

2.6

2.8

3.0

400

3.2

3.4

eV

nm

Optical Spectroscopy
Name: _________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
EXPERIMENT 7 Optical Spectroscopy
Results/Observations:
Part A: Observations of Emission Spectra.
1. Fluorescent Light Spectrum:
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

2. Incandescent Light bulb Spectrum:

1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

3. Helium Spectrum:
1.7

1.8

1.9

700

2.0

2.2

600

2.4

500

2.6

2.8

3.0

400

3.2

3.4

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

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Optical Spectroscopy
Name: ________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
4. Spectra of Salt Solutions:
a. Sodium Chloride Spectrum:
Observations (i.e. color of light, wavelengths from spectrometer):
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

b. Lithium Chloride Spectrum:

Observations (i.e. color of light, wavelengths from spectrometer):
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

c. Potassium Chloride Spectrum:

Observations (i.e. color of light, wavelengths from spectrometer):
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

d. Calcium Chloride Spectrum:

Observations (i.e. color of light, wavelengths from spectrometer):
1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

e. Strontium Chloride Spectrum:

Observations (i.e. color of light, wavelengths from spectrometer):
1.7

1.8

1.9

700

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2.0

2.2

600

2.4

500

2.6

2.8

3.0

400

3.2

3.4

eV

nm

Optical Spectroscopy
Name: ________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
Part B: Analysis of Spectra.
1. Test of Procedure for Determining the Composition of an Unknown Salt Solution:
Test mixture contains: _______________

1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

Unknown Salt Solution Spectrum and Identification Number:

1.7

1.8

1.9

2.0

700

2.2

600

2.4

2.6

2.8

500

3.0

3.2

3.4

400

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

The unknown salt solution contains: _________________________________

2. Hydrogen Spectrum:

1.7

1.8

1.9

700

2.0

2.2

600

2.4

500

2.6

2.8

3.0

400

3.2

3.4

eV

nm

Observations (i.e. color of light, wavelengths from spectrometer):

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Optical Spectroscopy
Name: ________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
Use the observed wavelengths from the hydrogen emission spectrum to calculate the
differences between hydrogen energy levels. (Show one sample calculation. Tabulate the rest
of the values.)

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Optical Spectroscopy
Name: ________________________________ Date: _________________________________
Lab Instructor: __________________________ Lab Section: ___________________________
Draw a partial energy-level diagram for hydrogen. Assume that all observed transitions
terminate at the n = 2 state; for example, if you observe two transitions, they are from state A
n = 2 and state B n = 2. Also, set the value of the energy of n = 2 state to 0.545 aJ.

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