EMISSION & ABSORPTION SPECTRA
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 with 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, this configuration is unstable. The "excited" electron
releases its newly acquired energy as it falls back to its initial or ground
state. Often, the excited electrons acquire sufficient energy to make
several energy level transitions. When these electrons return to the
ground state, several distinct energy emissions occur. The energy that
the electrons absorb is often of a thermal or electrical nature, and the energy
that an electron emits when returning to the ground state is called electromagnetic
radiation.
In 1900, Max Planck
studied visible emissions from hot glowing solids. He proposed that light
was emitted in "packets" 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 (n)
of emitted light and Planck's constant (h). The equation is written
as E = hn
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 spectra or emission 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 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.