8 —
Atomic Spectrometry
Electronic energy levels are widely spaced, and spectral transitions between them are therefore
observed towards the high energy end of the electromagnetic spectrum, i.e. in the visible, ultraviolet and
X-ray regions. The spectra are characterized by narrow lines and may be simple or complex depending
on the number of excited states involved in the excitation process. The wavelengths of the observed
absorption and emission lines are characteristic of a particular element, and the intensity of a given
spectral line is proportional to the number of atoms undergoing the corresponding transition. Atomic
spectrometric techniques thus provide means both for the qualitative identification of elements and for
their quantitative determination. Because of the large dipole changes associated with electronic
transitions, the sensitivity of the techniques is high and they are used primarily in the field of trace and
minor component elemental analysis. A summary of those used in analysis is given in Table 8.1. To
understand these techniques, it is helpful to review certain aspects of the theory of atomic structure and
electronic transitions.
Atomic Structure and Spectra
According to quantum theory, the electrons of an atom occupy quantized energy levels or orbitals
defined by a set of four quantum numbers whose permitted values (Table 8.2) are determined by
mathematical rules. The natural tendency is for electrons to occupy orbitals of the lowest possible
energy consistent with the Pauli exclusion principle which states that no two electrons in one atom may
be defined by the same set of values for the four quantum numbers n, l, ml and s. The shapes and
energies of the orbitals are determined by the values of the quantum numbers and by complex inter-
electronic effects. For most elements, therefore, orbital energies cannot be calculated exactly although
an energy level diagram showing relative values can be drawn. For example, part of the energy level
diagram for the outer orbitals of sodium is shown in Figure 8.1. In its ground state, the single valence
electron of sodium occupies the orbital defined by n = 3, l = 0, m = 0, s = –1/2. By suitable excitation,
the electron can be promoted to an orbital of higher energy from which it can return to the ground state
with emission