Atomic Emission Spectroscopy
Introduction
Spectroscopy is the study of the
interaction of light with matter.
Elements and compounds when exposed to large amounts of energy emit
light at certain frequencies. It is
possible to induce hydrogen atoms to emit visible light by exciting them, as by
applying a high electrical potential. This set of visible colors or lines is
called the Balmer series. Using an
optic arrangement called a spectroscope, it is possible to resolve the visible
light emitted into its component colors and to determine their wavelengths.
These wavelengths are characteristic of hydrogen and the energy emitted or
absorbed when the electron jumps from one energy level to another.
The energy levels of the hydrogen atom are
schematically represented in the following diagram:
ELECTRONIC
ENERGY LEVELS OF THE HYDROGEN ATOM
E
= 2.18 x 10-18J _______________________________n
= 
_______________________________n
= 6
_______________________________n
= 5
_______________________________n = 4
Energy
(Joules/photon) _______________________________ n = 3
_______________________________ n = 2
Ground State ________________________________n =
1
USEFUL PHYSICAL
CONSTANTS AND CONVERSION FACTORS
= 6.63 × 10-34 Joule-seconds.
c = 3.00 × 108 meter/seconds.
1 meter º 1 × 1010 Angstroms.
The energy of a
transition between levels can be calculated from:
|
|
(1) |
where
"outer" refers to outer electron energy level and "inner"
refers to inner electron energy level.
By using Bohr's postulates it can be
shown that for the hydrogen atom
|
|
(2) |
where n = 1, 2, 3,.... and refers to
an allowed energy level. (See above diagram) Thus, by measuring the wavelength
of the lines in the hydrogen spectrum, we can use the above formula to find the
initial and final energy levels of the electron. All the transitions in the
visible spectrum of hydrogen terminate on the n = 2 level.
Just as hydrogen has its
characteristic spectrum, so do all the other elements. In fact the spectrum of
the gaseous atoms of an element can be used to identify it. In the second part
of this experiment you will be given several lamps containing
"unknown" gaseous atoms. You will identify these unknowns by
measuring the wavelengths of the lines in their spectra and comparing them with
the tabulated spectral lines. The Bohr
equation can only be used for hydrogen.
Equipment
Spectroscopes,
5000 volt transformer, lamps containing H, He, Hg, and Ne.
Procedure
DO
NOT TOUCH THE LAMP OR THE METAL CONNECTIONS WHILE THE APPARATUS IS ON
1.
Insert the hydrogen lamp into the 5000 volt source and turn it on. Only touch
the ends of the lamp. Never touch the
middle. Line up the slit in the
spectroscope with the lamp and observe the light.
2.
As you look into the spectroscope, look to the right and you will see a scale superimposed on the
spectral lines. Record the wavelengths of the lines observed to 2 significant
figures. The scale is in angstrom units. One angstrom unit equals 1 x 10-10 meters.
The units are different for the hand held spectroscope. You should be able to see three or four
lines. Use this data to calculate the
initial level of the electron for each transition (no in equation
1).
3.
Use this same procedure for each of the 3 "unknown" lamps. For some of the unknowns, you will see an
almost continuous spectrum rather than discrete lines. Record the range and center wavelength of
this broad band of color.
Compare the
wavelengths to those in Table 1 and determine the identity of each unknown.