Geoscience Reference
In-Depth Information
bonding interactions at all. In many systems molecules may exhibit molecular orbitals that
change state from antibonding to bonding or vice versa. This is entirely dependent on the
atoms/conjugated system involved, and the nature of the bonding is relative to the atoms
concerned. In more conjugated systems involving molecules with several atoms, for exam-
ple, benzene, a particular molecular orbital may be bonding with respect to some adjacent
pairs of atoms and antibonding with respect to other pairs. In this instance the ratio of
bonding to antibonding molecular orbitals becomes important. For example, if the number
of
bonding interactions
outnumbers the
antibonding interactions
, the molecular orbital in
question is deemed to be
bonding
and vice versa. For benzene each carbon atom contrib-
utes only one electron to the delocalized
π
-system of benzene, and because there are only
six
π
-electrons only the three lowest-energy (bonding) molecular orbitals are filled.
1.2.2.4 Nonbonded Electrons
Valence electrons that are not used for bonding must be paired and are known as “lone pair
electrons” or “nonbonding electrons” (
n
electrons). Almost all atoms have paired electrons
in the valence shell. Although much attention is given to unpaired electrons in atoms, non-
bonding electrons are important in determining the geometry of a molecule along with
unpaired electrons using the VSEPR theory. Lone pair electrons exhibit higher energies
than
σ
or
π
pairs because their repulsive forces are greater. In a
σ
bond a bonding electron
pair lies farther away from the central atom than does a lone pair. Therefore, if the overall
geometry of a molecule has two sets of possible positions, but each position has a differ-
ent degree of repulsion, then lone pairs will occupy the position that has less repulsion.
Although paired electrons are not specifically involved, in bonding interactions, they can
contribute to the spectral features of a molecule and therefore must be considered.
1.3 Understanding the Fluorescence Process
British scientist Sir George G. Stokes first described “fluorescence” in
1852
after the blue-
white fluorescent mineral fluorite (fluorspar). Stokes is perhaps better remembered for dis-
covery of the observed differences (in wavelength or frequency) in positions of the band
maxima of the absorption and emission spectra of the same electronic transition, the so-
called Stokes Shift. There are three fundamental processes that are implicit in the emission
of light by a molecule (Lakowitz,
2006
):
excitation
of the molecule (i.e., absorption of an
appropriate photon),
vibrational relaxation
(nonradiative decay), and finally the
emission
of light
(radiative decay). All of these processes occur on timescales that are separated
by several orders of magnitude (see
Table 1.2
). Excitation of a molecule by an incoming
photon happens instantaneously (femtoseconds, 10E
-15
), while the vibrational relaxation
of an electron in an excited state to the lowest energy level is slower, usually occurring
over picoseconds (10
-12
). The final process, light emission, almost always occurs at longer
wavelengths and the return of the molecule back to the ground state occurs over nanosec-
onds (10
-9
seconds).
