BACKGROUND: Spectroscopy is the study of the manner in
which light (electromagnetic radiation) interacts with materials. As a
scientific discipline, spectroscopy occupies that fertile ground between
chemistry and physics. Fundamental investigations in spectroscopy have led to
our understandings of the electronic structures of atoms and the geometry and
bonding of molecules. Spectroscopy research today focuses on the fundamental
aspects of the energetics and mechanisms of chemical
reactivity. Applied spectroscopy exploits the fact that the wavelengths
(energies) of light that is absorbed or emitted by a material are unique to
that material. In various modifications, this principle can allow either
qualitative identification of the components in a sample or quantitative
determination of the amounts of an element or compound in a sample. In this
exercise we will use concepts from both fundamental and applied spectroscopy.
Convenient for the purposes of this
experiment, the electronic spectra of the elements fall largely within the
range of energies (wavelengths) that are detected by the human eye. In Part A,
we will work with five Group I & II elements and use subjective visual sense to determine the colors emitted by these elements, and also
use the student spectroscope to determine the specific wavelengths of
the spectral lines emitted by these elements. From this information
we can determine the identity of the cation in a solution containing a single
metal halide.
One of the curious observations of
the mid-nineteenth century that defied explanation for nearly fifty years was
the fact that unlike the continuous "rainbow" spectra emitted by the
sun or a white light, the spectra emitted by a single element consisted of a
small number of discrete lines. Finally in 1914, Niels
Bohr provided an explanation for this phenomenon. Bohr's quantum mechanical
theory suggested that atoms of an element emit radiation when they undergo
transitions between two energy states. And further, the two energy states of
the atom differ with respect to the orbital radius of the electrons. The theory
provides a method to calculate the energy of these states
En
= (-2.18 x 10-18 J) ( Z2 / n2)
,
where Z is the number of protons in the nucleus of atoms of that
element, and n is a positive integer identifying the energy state of the atom.
Using the Planck relationship ( ΔE
= h c / l ) and an expression for change in energy (ΔE = Efinal
- Einitial) scientists were able to
confirm the essential validity of the Bohr quantum mechanical model.
In Part B, we will calibrate the
wavelength scale of the student spectrometer using the known wavelengths of the
helium emission lines. Then using that calibration, students will determine the
wavelength of the most prominent lines of hydrogen and calculate the atomic
state transition responsible for those hydrogen lines.
Plank's
constant, h = 6.63 x 10-34 J-s
Speed of light, c = 3.00 x 108 m/s
Note: The unknown
samples used in this course contain only a single metal cation. Qualitative
analysis of more complex samples containing two, three or even many elements
can be accomplished by similar, though more complicated, instrumental
techniques. Those techniques are studied and used in later courses.
|
Color of
helium emission line |
Accepted Wavelength |
|
red |
668 nm |
|
yellow |
588 nm |
|
green |
502 nm |
|
green |
492 nm |
|
blue-green |
471 nm |
|
blue-violet |
447 nm |
|
violet |
403 nm |