Transcript Tying some loose ends and introducing some new ones.

Valence Bond Theory
SCH4U1
September 22 2011
Mr. Dvorsky
Valence Bond Theory
• Valence Bond Theory tries to tackle this
problem.
First, the some main principles of the theory.
There are three:
1. Covalent bonds between two nuclei are
formed when each atom contributes one
valence electron to a common orbital. This
common orbital, which is an overlap of atomic
orbitals, contains two electrons of opposite
spin.
• For example, in H2, the overlap of the
spherical 1s orbital from each H atom results
in the formation of a sigma bond.
-a sigma bond is defined as a bond formed by
overlap of orbitals in a region of space that
lies on the same axis as the two nuclei (headon overlap).
-the sigma bond is cylindrically symmetrical.
2. As a result of the overlap, the electrons are
deemed to be localized. In other words, they
are restricted to the area between the two
respective nuclei. They do not move around
throughout the molecule.
3. Within the molecule, the atomic orbitals
located on the central atom are not
necessarily “pure” atomic orbitals. Bonds
involving elements in the second or higher
row in the periodic table involve combinations
of atomic orbitals that form “hybrid” orbitals
as the bonds are being made.
• Each sp3 hybrid orbital on the carbon contains
one electrons, and each hybrid orbital can
then directly overlap with the 1s orbital of
hydrogen to form a sigma bond.
-after this overlap, there are two electrons
between the nuclei for each bond created.
-109.5 degrees between each hybrid orbital (like
a tetrahedron).
Hybrid orbitals can also contain non-bonding
electrons. Ammonia, NH3, the lone pair of
electrons is also in a hybrid sp3 orbital.
-lone pairs have an effect
nn shape which we will
Get into with VSEPR
It is important to realize that sp3 hybridization is
just one possibility. Our course will include 5
major hybridization types:
Atomic Orbitals Used
one s, one p
one s, two p
one s, three p
one s, three p, one d
one s, three p, two d
Hybrid Orbitals Formed
two sp
three sp2
four sp3
five sp3d
six sp3d2
sp Hybridization
-the combination of one s and one p results in
the formation of two sp orbitals. These two
orbitals are 180 degrees apart, the optimum
angle that separates two regions of electron
density.
-With only only of the p orbitals used, the two
remaining p orbitals are perpendicular to each
other and to the sp hybrids. –can be used later
to make double and triple bonds as we will
discuss in a few moments.
sp2 Hybridization
-in this type of hybridization, one s and two p
orbitals combine to form three sp2 hybrids that
are 120 degrees apart. The remaining p orbital
on each carbon atom is perpendicular to this
plane and may be used to make a double bond.
sp3 hybridization
-As shown previously in our CH4 example, one s
and all three p orbitals combine to form four sp3
hybrids in a tetrahedral arrangement. After
combining, there are no remaining p orbitals so
atoms that are sp3 hybridized cannot form
double or triple bonds.
sp3d and sp3d2 hybridization
-these types are not possible with elements in
the second row (C,N,O, etc) where valence
electrons are in n=2 (i.e. there are no d orbitals).
With elements in the third row or higher, d
orbitals are available so these common
hybridization types are possible:
One s + three p + one d = five sp3d orbitals (four d
orbitals remain unused)
One s + three p + two d = six sp3d2 (three d orbitals
remain unused)
Lewis structures fail to explain how things like
PCl5 are possible