The Principles of Fluorescence
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than the 1σ orbital, and also higher than either of the individual atomic orbitals. This can be
explained in terms of destructive interference of the two orbitals. The point in space which
is equidistant from each nuclei and intersects the internuclear axis has a wavefunction
equal to zero. This is known as the nodal plane and is referred to briefly in Section 1.2.2.
Both opposing orbitals cancel each other on this plane as result of destructive interference.
This 2σ bonding is known as an antibonding orbital and is denoted as 2σ *. This orbital
excludes the electron from the internuclear region and relocates it to outside the region of
bonding. The net result of this electron relocation is that the orbital is repulsive, pulling the
nuclei apart. This is the main reason why the antibonding molecular orbital 2σ * exhibits
higher energy than the 1σ molecular orbital.
A molecular orbital exhibits “antibonding” properties when the electron density between
the two nuclei concerned is lower than would otherwise be predicted if there were no
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 example, 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
contributes 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, nonbonding 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 different 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 bluewhite fluorescent mineral fluorite (fluorspar). Stokes is perhaps better remembered for discovery of the observed differences (in wavelength or frequency) in positions of the band
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than the 1σ orbital, and also higher than either of the individual atomic orbitals. This can be
explained in terms of destructive interference of the two orbitals. The point in space which
is equidistant from each nuclei and intersects the internuclear axis has a wavefunction
equal to zero. This is known as the nodal plane and is referred to briefly in Section 1.2.2.
Both opposing orbitals cancel each other on this plane as result of destructive interference.
This 2σ bonding is known as an antibonding orbital and is denoted as 2σ *. This orbital
excludes the electron from the internuclear region and relocates it to outside the region of
bonding. The net result of this electron relocation is that the orbital is repulsive, pulling the
nuclei apart. This is the main reason why the antibonding molecular orbital 2σ * exhibits
higher energy than the 1σ molecular orbital.
A molecular orbital exhibits “antibonding” properties when the electron density between
the two nuclei concerned is lower than would otherwise be predicted if there were no
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 example, 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
contributes 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, nonbonding 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 different 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 bluewhite fluorescent mineral fluorite (fluorspar). Stokes is perhaps better remembered for discovery of the observed differences (in wavelength or frequency) in positions of the band