UV-Visible and Infrared Spectroscopic Methods in Environmental Analysis - Fundamentals of Environmental Sampling and Analysis - page 195

y y y

y

x x x x

z z z

z

Figure 8.5 Electron in s and p orbital

2s orbital 2p x orbital 2p y orbital 2p z orbital

Note that the closer the orbital to the nucleus the lower is the electron's energy. The
relative energies of atomic orbitals are: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s <
4d

< 5p < 6s < 4f < 5d < 6p < 7s < 5f.

Types of Absorbing Electrons: r, r*, p, p* and n

We now look at two major types of orbitals (s and p) and how electrons in these
atomic orbitals form molecular bonding. The s electron orbital has a spherical shape
that is also symmetrical without directional characteristics, whereas p electron
orbital is dumbbell shaped. Since there are three p atomic orbitals (Table 8.1) and
they have equal energy, the three orbitals are perpendicular to each other,
symmetrical about the x-, y-, and z- axis and are designated as 2p x ,2p y , and 2p z
orbitals to emphasize their directional nature (Fig. 8.5).
When two s atomic orbitals overlap (such as two 1s hydrogen atoms combine
to form hydrogen molecule H 2 ), the bond that is formed is called a sigma (s) bond
(Fig. 8.6, bottom). During bond formation, energy is released as the two orbitals
start to overlap till maximum stability (minimum energy) is achieved when two
nuclei are at a certain distance apart. This distance corresponds to the bond
length , and the energy released corresponds to the bond energy . To break the
bond, the same amount of energy is required (termed the bond dissociation
energy).
There is another way to combine 1s atomic orbitals termed sigma antibonding
molecular orbital (s*) (Fig. 8.6, top). An antibonding orbital can be perceived as
the detraction from formation of bond between two atoms, similar to the
darkness that occurs when two light waves cancel each other or the silence that
occurs when two sound waves cancel each other.
In addition to the formation from two s orbitals as described in Figure 8.6, the
s orbitals can also be formed when two 2p atomic orbitals overlap end-on (i.e., head
to head), such as two 2p orbitals of fluorine (1s 2 ,2s 2 ,2p 5 ) overlap to form fluorine
gas (F 2 ).

σ∗ antiboding molecular orbital

Figure 8.6 The formation of sigma
(s) bonding molecular orbital and s*
antibonding molecular orbital. Prior to the
formation of a covalent bond, each
electron is in 1s atomic orbital. These two
electrons are in the molecular orbital
after the formation of a covalent bond

1s atomic
orbital

1s atomic
orbital

+

σ boding molecular orbital