operator are usually given in the character table, as discussed in the previous
chapter regarding allowed vibrational transitions.
For polyatomic molecules, the point group has enough symmetry elements
(or rather, classes, and so therefore irreducible representations) that the fol-
lowing statement is usually applicable: the ground electronic state and the al-
lowed excited states are usuallyof different irreducible representation labels.
There are some general rules for electronic spectra of molecules (although
exceptions to such rules are not uncommon). For most molecules that are
composed of atoms of main-group elements that have an electron configura-
tion with a saturated valence shell, most of the low-energy electronic transi-
tions are already relatively high in energy and already require UV light (that is,
higher energy than visible light) for an allowed transition. This is why mole-
cules such as water, ammonia, methane, and so on are colorless. They do not
absorb visible light because the electronic transitions are caused by invisible
UV light.
In molecules that have an atom with an unpaired electron, there is a good
chance that relatively low-energy visible light is energetic enough to cause an
electronic transition. An example is NO 2 , a rare case of a stable main-group
compound that has an odd number of electrons. It is brown, and is largely re-
sponsible for the color of smog. This idea is particularly applicable to com-
pounds that contain d-block or f-block elements: transition, lanthanide, and
actinide atoms. Consider compounds that have a Cu^2 ion in them. Such an
ion has the valence shell electron configuration 3d^9. There is a single unpaired
electron. Therefore, one would predict that Cu^2 compounds are colored, and
they usually are. However, consider Zn^2. It has a valence shell electron con-
figuration of 3d^10 , having no unpaired electrons. Zinc compounds are not
known for their colors. (Yes, it is understood that these examples are ions, not
molecules. They are simple examples, and since such cations are never present
without an anion, we are not stretching the definition toomuch.)Example 15.12
Predict whether the following molecules would have color. That is, will elec-
tronic transitions occur in the visible region of the spectrum, or will they prob-
ably occur in the invisible UV region of the spectrum? State the reason(s) why.
a.Sodium chloride, NaCl
b.Iron pentacarbonyl, Fe(CO) 5 (Consider the ligands and the metal atom
separately.)
c.Chloroform, CHCl 3
d.Titanium dioxide, TiO 2
e.Hemoglobin, which has four iron atoms in itSolution
a.Both the sodium ion, Na^ , and the chloride ion, Cl^ , have an octet elec-
tron configuration for the valence shell. All the electrons are paired, so one
would expect that sodium chloride would not absorb in the visible region of
the spectrum. Crystalline NaCl is colorless and can be used for optical com-
ponents.
b.Although the ligands have all-paired electrons, the neutral iron atom has a
3 d^6 electron configuration, so one might predict that iron pentacarbonyl
would absorb in the visible region of the spectrum. Iron pentacarbonyl is a
volatile liquid at room temperature and has a yellow color.542 CHAPTER 15 Introduction to Electronic Spectroscopy and Structure