Electron Transitions in Optical Emission and X-Ray Fluorescence Spectrometry - - Spectroscopy
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Electron Transitions in Optical Emission and X-Ray Fluorescence Spectrometry


Spectroscopy



Volker Thomsen
The origin of the characteristic spectral lines of the elements in the optical and X-ray regions of the electromagnetic (EM) spectrum is all about electron transitions. Excitation of outer-shell electrons to an excited state and their subsequent return to ground state results in electromagnetic radiation in the optical (ultraviolet–visible–infrared) region with wavelengths from about 100 to 1000 nm. Excitation of inner-shell electrons produces EM radiation in the X-ray region with wavelengths from about 0.01 to 10 nm.


Table I: Prominent spectral lines of magnesium
Figure 1 shows the two X-ray spectral lines and some prominent optical emission lines of magnesium (Z = 12). Note the logarithmic scale for wavelength. These spectral lines are listed by both wavelength and energy in Table I. Some relative spectral line intensities are also shown (2). The designation "U" for the optical emission lines refers to spectral lines of the neutral atom and "V" for the singly-ionized atom. The number that follows indicates the relative intensity with "1" generally being the strongest spectral line, depending upon the excitation condition.


Figure 1: Spectral lines of magnesium in the EM spectrum.
Some additional comments on Figure 1 and its associated table:

  • The "U" and "V" nomenclatures are somewhat antiquated, included here solely to preserve the connection to the source, the MIT Wavelength Tables (2). Currently, spectral lines in the optical region are designated as follows: A capital "i" (I) for the neutral atom, II for the first ionized state, and so forth. For example, today we write Mg I 285.21 nm and Mg II 279.55 nm in referring to these spectral lines.
  • Please note that the X-ray and optical relative intensities of Table I are on different scales (3).

Electron Energy Levels


Figure 2: Generic energy level diagram.
The electron energy levels of atoms are grouped together based upon the principal quantum number n. These groups differ from each other by about a factor of 10 in binding energy. Figure 2 shows a generic electron energy level diagram for the first three levels (n = 1, 2, 3) and the associated sublevels. The standard electron energy level designations are shown along the left side along with those used in X-ray fluorescence. On the right side are shown the associated quantum numbers: the principal quantum number "n," the orbital angular momentum quantum number, "l," and the magnetic quantum number "m." This diagram sets the stage for those to follow.

Electron Energy Levels: X-ray

The electron energy transitions resulting in X-rays are easier to portray. The innermost electron shell is called the K-shell; the next farthest out from the nucleus is called the L-shell, followed by the M-shell, and so forth.




To produce the characteristic spectral lines in the X-ray region, the excitation energy (E) must be such that it is equal to or greater than the energy required to remove an electron from its shell, known as the "binding energy" (EB). That is, we must have:


Figure 3: X-ray fluorescence process.
An outer shell electron will "drop down" to fill the void created in the inner shell and there is a certain probability that an X-ray characteristic of that atom will be emitted. This process is called X-ray fluorescence (XRF) and is shown in Figure 3 (4). The characteristic X-rays are called "K lines" if they result from an electron filling the innermost, or K-shell, and "L lines" if they result from filling the next electron shell out, the L-shell. Figure 4 shows an X-ray energy level diagram for magnesium.


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