Name:
Lab Partner:
Date:
Emission Spectra and Energy Levels
OBJECTIVES
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Observe visible emission spectra of several elements with the aid of a simple
hand-held spectrometer.
Connect the emission spectra to the electron configuration of elements.
Demonstrate emission spectra analysis as a powerful analytical tool.
INTRODUCTION
Emission spectra are produced when excited electrons in atoms return to a more stable state,
called the ground state, by emitting radiation of specific wavelengths. One convenient method
of exciting the atoms of an element is to pass an electric current through a sample of the
element in the vapor state. This is the principle behind the spectrum tubes you will use in this
experiment. A spectrum tube contains a small sample of an element in the gaseous state. An
electric current through the tube will cause the vapor to glow brightly. The light (glow) is
produced when excited electrons emit energy in the visible light wavelengths as they return to
their ground states.
When visible radiant energy from a spectrum tube is passed through a diffraction grating in the
spectrometer, an emission spectrum is produced. Each element has its own unique electron
configuration and, therefore, its own unique emission spectrum. Such a spectrum consists of a
series of distinct bright lines of definite wavelength. Each wavelength is related to a specific
quantity of energy produced by an electron jumping from its excited state to a lower energy
level. Therefore, emission spectra are experimental proof that electrons exist at definite,
distinct energy levels in an atom.
This diagram represents the Bohr model of the hydrogen atom and shows how the visible
spectrum is generated by the excited electrons inside the spectrum tube. In this experiment,
you will study the emission spectra of several elements.
Procedure
Record your observations for each element on the blank spectrum sheets
provided.
1. Identify the element in the spectrum tube that has been inserted into the high voltage
transformer.
2. Turn on the transformer. Use extreme caution because the spectrum tubes become
very hot.
3. Look through one of the hand-held spectrascopes aimed directly at the glowing
spectrum tube. Line up the slit in the spectrascope with the glowing spectrum tube.
4. Record the bright spectral lines on the blank spectrum sheets using colored pencils,
matching the colors and locations (wavelengths) as accurately as possible.
5. Repeat steps 1 through 4 until all of the spectrum tubes have been viewed.
6. When the experiment is concluded, turn off the transformer but do not remove the
spectrum tube. The teacher will remove the spectrum tubes since they are easily
broken.
DATA and OBSERVATIONS
Record your data on the blank emission spectra graphs attached.
ANALYSIS and CONCLUSIONS
Questions:
1. Explain why elements larger than hydrogen have more complicated emission spectra.
2. How can emission spectra be used to study the composition of stars?
3. In 1900, German physicist Max Planck concluded that matter can gain or lose energy in
small, specific amounts called quanta. A quantum is the minimum amount of energy that can
be gained or lost by an atom. Planck’s Law states that the energy emitted by a substance was
directly related to the frequency of that energy:
E quantum = hν (h = Planck’s constant = 6.626 x 10-34J.s)
Frequency, ν, is inversely related to wavelength,λ, by the following equation:
c = λν (c = speed of light = 3.0 x 108 m/s)
In the table below, calculate the energy emitted by the visible light for hydrogen. Multiply the
number in the graph by 100 to get the nm value.
The Emission Spectrum for Hydrogen
Color
Wavelength (nm)
Frequency (1/s)
Energy (J)
Red
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Blue
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Violet
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Violet (faint) ____________
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