Earlier the information from line spectra was used in the development of the modern Quantum Mechanical model of the atom. The hydrogen spectrum is frequently used in describing the electron transitions. Since hydrogen has only one electron, all the orbitals at any energy level (s, p, d, f) are degenerate (the same energy). When the spectra of poly-electron atoms are examined the results can be much more complex than hydrogen since the orbitals that are occupied are not degenerate. This results in more lines and in fainter lines in many cases.
Review the basics of electromagnetic emission and absorption by atoms.
As electron transitions are discussed they generally fit into one of four areas of the electromagnetic spectrum. Relatively small transitions of electrons, such as from the 4th energy level of hydrogen to the 3rd (and all smaller transitions), have energies that put them in the infrared portion of the spectrum. Larger transitions involving energies between 4.9x10-19 Joules and 2.8x10-19 Joules lie in the visible region of the spectrum. The transitions in hydrogen that correspond to this range are from the 2nd energy level to the 3rd, 4th, 5th, and 6th energy levels. In other atoms they may involve similar transitions; generally elements with more protons will have bigger jumps between the first three levels than does hydrogen.
In hydrogen a transition from any higher level to the first energy level will fall in the ultraviolet portion of the spectrum. In other atoms, transitions involving the first, second, and third level may fall in the UV range.
Since elements with a greater nuclear charge than hydrogen hold the first levels of electrons more tightly, more energy is required to promote those electrons. In these cases, the transitions to or from the first and second levels usually fits in the region of the electromagnetic spectrum that we call x-rays. Recall that Henry Moseley used the pattern of x-rays generated by bombarding different metals with high energy to establish the charge on the nucleus.
