Geoscience Reference
In-Depth Information
that this is not the case. Therefore the energy orbitals within atoms (
s, p
,
d
, and
f
) also pos-
sess sublevels (principle quantum number) of energy. Owing to the wave-particle duality
of electrons it is impossible to pinpoint their exact position; instead we can only consider
the probability of finding an electron in a region of space. The probability (expressed as the
volume around a nucleus in which an electron is 90% probable to be found) is referred to as
the atomic orbital. The orbital of
s
electrons is spherical (
s
orbital). There are three different
p
orbitals (
p
x
,
p
y
,
and
p
z
), which have equal energies but different directions in space. These
are often described as dumbbell orbitals. For electrons at higher energies,
d
and
f
orbits
become available. These orbitals are more complex and numerous than the orbits observed
for
s
and
p
electrons.
The distribution of electronic charge throughout the axis of a chemical bond is impor-
tant. In covalent bonding the region where the value of an orbital wave function (
ψ
) equals
zero (or is very low),defines a region of space within the system where there is zero elec-
tron density. This is known as a nodal plane, and quantum theory indicates that molecu-
lar orbitals with identical symmetries mix, and the wave functions for
s
+
s
and
p
z
+
p
z
become blended. The extent of this mixing (or blending) depends on the relative energies
of the molecular orbitals involved and is extremely important for determining the number
of nodal planes and distribution of energy within molecular bonding orbitals. This mixing
of wave functions is termed resonance. Typically, molecules exist as a number of atoms
bonded together via covalent bonding, and the collective arrangement of these atoms is
such that the overall molecular structure is electrically neutral. Within this structure, all
outermost electrons of the atoms involved are paired with other electrons, either in bonds
or lone pairs. These outer electrons are termed
valence electrons
and are very influential in
determining how atoms interact with each other (reactivity).
Lewis's original theory could not take into account the shape adopted by molecules.
Gillespie and Nyholm (
1957
) developed the currently accepted modern theory of chemical
bond formation (MO and VB theories), which uses the valence-shell electron pair repulsion
model (VSEPR) to account for molecular structure (Gillespie,
1970
). VSEPR states that
molecular shape is caused by repulsions between electron pairs in the valence shell.
1.2.2.1 Sigma Bonds (σ Bonds)
Sigma bonds (
σ
bonds) are the strongest type of covalent chemical bond and are perhaps
best illustrated in simple diatomic molecules such as H
2
, F
2
, Cl
2
, Br
2
, and I
2
. Sigma bonding
in diatomic molecules is always symmetrical with respect to the rotation about the bond
axis (nucleus to nucleus). Therefore common
σ
bonds can be represented as
s
+
s
,
p
z
+
p
z
,
s
+
p
z
,
and
d
z
2
+
d
z
2
(where
z
is defined as the axis of the bond). In
σ
covalent bonding the
two “shared” electrons can either originate from the same atom, in which case the
σ
bond
is covalent, or from each atom, where the
σ
bond is termed a coordinate covalent bond. For
homo diatomic molecules, bonding
σ
orbitals have no nodal planes between the bonded
atoms, whereas in the case of hetero diatomic atoms forming a covalent bond (where one
atom is more electronegative than the other) the electron pair will spend more time closer
to that atom. This is termed a polar covalent bond.
