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THE HYDROGEN-ATOM SPECTRUM
Another piece of strange behavior observed at about the same time is the so-called "line spectrum" of the hydrogen atom. You are already familiar with the fact that a prism can be used to disperse (separate) the colors of white light (like light from the sun or a light bulb) into its component colors: Red, Orange, Yellow, Green, Blue, Indigo and Violet ( ROYGBIV ). The rainbow that results from dispersion by a prism is called the spectrum of the sun.
We know that our sun is made up of many elements, everything, in fact, from hydrogen (most abundant) to Iron (0.003% of all atoms in the sun). Although sunlight consists of more of some colors and less of others, we refer to its light as white light because it contains all colors visible to humans. However, if we look at the emission spectrum of only hot, glowing hydrogen gas (e.g. by putting H 2 gas in a fluorescent bulb), dispersing it with a prism in just the same way, we see a completely different result (figure below). Now instead of a rainbow, we see only four distinct "lines" or "bands" of colored light with nothing in between - thus the term "line spectrum."
The two spectra below show absorption by and emission from an ensemble of hydrogen atoms.
In the top panel, we see that when white light is shined on a sample of hydrogen gas, the H absorbs certain discrete wavelengths of light, leaving the rest to pass through the prism and be dispersed into colors. The black bands represent missing light - light that has been absorbed by the hydrogen.
In the lower panel, Hydrogen gas has been excited with electricity in a kind of fluorescent tube, and passed through the prism. The light emanating from the excited hydrogen atoms consists of only four discrete color bands, red, cyan, blue and violet. The wavelengths of the colors are given (in nanometers), and form a characteristic fingerprint of Hydrogen.
The actual emission spectrum of the sun (and anything else) is also really a line spectrum, but because there are so many different types of atoms in the sun, and because all have much more dense and complex line spectra than H(which is the simplest atom) they overlap and blur together to give our familiar rainbow.
We assume that electricity excites the electron in an H atom to one of the " excited states " - higher energy levels than the ground state or lowest energy level.
When those excited electrons relax , they emit light that exactly equals the difference between the two energy levels that define the transition.
This must be so because energy must be conserved — neither created out of nothing, nor destroyed. You can see that the visible line spectrum represents transitions from excited states to the second-highest energy level. The spectra of larger atoms are proportionally more complicated.
The absorption spectrum is just the reverse. The electron in a ground-state hydrogen atom absorbs certain energies of light in order to achieve an excited state.
The idea that the energies of electrons in atoms are discrete and not continuous is another example of the non-classical (not like objects we're used to seeing) and unexpected behavior of electrons that begged for a non-classical explanation.
Source: http://www.drcruzan.com/Chemistry_Electrons.html
