Part A: EMISSION SPECTROSCOPY
According to the Bohr atomic model, electrons orbit the nucleus within specific energy levels.
These levels are defined by unique amounts of energy. Electrons possessing the lowest energy
are found in the levels closest to the nucleus. Electrons of higher energy are located in
progressively more distant energy levels.
If an electron absorbs sufficient energy to bridge the “gap” between energy levels, the
electron may jump to a higher level. Since this change results in a vacant lower orbital, the
configuration is unstable. The “excited” electron releases its newly acquired energy and falls
back to its initial or ground state. Often, the excited electrons acquire sufficient to make several
energy level transitions. When these electrons return to their ground state, several distinct
energy emissions occur. The energy that electrons absorb is often of a thermal or electrical
nature, and the energy that electrons emit when returning to ground state is electromagnetic
radiation.
In 1900, Max Planck studied visible emissions from hot glowing solids. He proposed
that light was emitted in “packet” of energy called quanta, and that the energy of each packet was
proportional to the frequency of the light wave. According to Einstein and Planck, the energy of
the packet could be expressed as the product of the frequency ( ) of emitted light and Planck’s
constant (h).
E=h
If white light passes through a prism or diffraction grating, its component wavelengths
are bent at different angles. This process produces a rainbow of distinct colors known as a
continuous spectrum. If, however, the light emitted from hot gases or energized ions is viewed
in a similar manner, isolated bands of color are observed. These bands form characteristic
patterns, unique to each element. They are known as bright line or emissions spectra.
By analyzing the emission spectrum of hydrogen gas, Bohr was able to calculate the
energy content of the major electron levels. Although the electron structure as suggested by his
planetary atomic model has been modified according to modern quantum theory, his description
and analysis of spectral emission lines are still valid.
In addition to the fundamental role of spectroscopy played in the development of today’s
atomic model, this technique can also be used in the identification of elements. Since the atoms
of each element contain unique arrangements of electrons, emission lines can be used as
“spectral fingerprints.” Even without a spectroscope, this type of identification is possible since
the major spectral lines will alter the color of the flame.
In the following experiment, you will use a spectroscope to examine several continuous
and bright line spectra.
OBJECTIVES:
1. To understand the relationship between atomic structure and emission spectroscopy.
2. To use a spectroscope and observe several sources of continuous and bright line spectra.
MATERIALS:
Colored pencils
Diffraction grating or spectroscope
Gas discharge tubes (TBA)
Power Supply
PRELAB QUESTIONS (Answer Questions in Lab Notebook)
1. According to Bohr’s atomic model, where may an atom’s electrons be found?
2. How do electrons become “excited?”
3. State the equation that is used to determine the energy content of a packet of light of specific
frequency.
4. What form of energy emission accompanies the return of excited electrons to the ground
state?
6. Why are students not permitted to touch the spectrum tube or high-voltage assembly?
PROCEDURE
1. Using the spectroscope, observe the continuous “rainbow” spectrum from an incandescent
light bulb.
2. Observe the colors of light in the visible spectrum and the wavelength range for each color
band. Sketch the spectrum of white light using colored pencils in the appropriate wavelength
boxes in the Spectrum Table. Note that the units of wavelength on the spectroscope are
nanometers (1 nm = 10-9 m)
3. Since the high voltages required for this part of the experiment may present an electrical
hazard, you should not touch the spectrum tube or high-voltage assembly. Your teacher will
load the spectrum tube and assemble the high-voltage power supply.
4. Once the power supply is turned on, aim the spectroscope at the glowing tube. Observe the
atomic emission spectrum of the gas. Work with a partner to note the principal features in the
spectrum.
5. Record the following information in the Data Table for the emission spectrum: the number of
lines, their colors, and their approximate wavelengths.
6. Using colored pencils, sketch the atomic spectrum of hydrogen in the wavelength boxes in
the Data Table.
7. Repeat for each of the gas samples and record your observations.
DATA AND OBSERVATIONS
Cut and tape table into lab notebook. Tape all around the edges (See Below) –Remember to
leave room to write in the table.
Table 1: (Remember to write and appropriate title)
Light Source
Incandescent
light
Spectrum
(number of
Lines)
700
650
Colors (Wavelength, nm)
600
550
500
450
400
CONCLUSIONS and QUESTIONS
1. How is a continuous spectrum different from an atomic emission spectrum? Explain.
2. According to Niels Bohr’s equation, the frequency (v) of light is inversely
proportional to its wavelength ( ) – as the wavelength increases, its energy
decreases. Based on the spectrum observed for incandescent white light, rank the
colors in the visible spectrum from highest energy to lowest energy.
3. Do all of the colors of light in the visible spectrum span about the same wavelength
“width” – that is, do the bands of color appear equally wide or narrow?
4. What color of light in the visible spectrum appears brightest? Does this mean that
it is the highest energy light?