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.