U.Mars — Encyclopedia
Basic Astronomy and the Nighttime Sky
Within an atom, small negatively-charged electrons orbit the nucleus (the latter which contains positively-charged protons, and neutrons). Rather than orbiting in completely-random configurations, the electrons in atoms occupy discrete energy levels, with the energies of those levels being different depending on the chemical element (its atomic number and mass).
Electrons can move between energy levels in atoms, but they must either gain or lose a very precise amount of energy corresponding to the energy difference between those levels. Recall that photons, or massless "particles" of light, carry energy; a photon of a particular energy has a particular wavelength (or frequency) related to that energy. Visible light photons have wavelengths in the range of about 400 to 800 nanometers. Therefore, certain "colors" of visible light, or photons of those colors, have just the right energy to match certain electrons' "jump" from one energy level in an atom to another (numerous photons in other regions of the overall electromagnetic spectrum also correspond to certain electron energy level differences).
For an electron to change from a lower energy state to a higher one, it must absorb a photon of the right energy. If an electron drops down from a higher-energy level to a lower one, a photon of that energy difference is produced and escapes. See Types of Spectra for more information on how these energy-level changes manifest in observed spectra.
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The higher the atomic number of an atom, the more electrons in it, and the more possible energy level trasitions there may be. Spectra for specific chemical elements can become very complicated, and difficult to distinguish from each other, but the specific array of energy levels in each element does serve as a kind of fingerprint for it.
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