1. SCH4U1 September 22 2011 Mr. Dvorsky
Valence Bond Theory
SCH4U1
September
Mr. Dvorsky

6. 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.

7. 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 (head-on overlap).
-the sigma bond is cylindrically symmetrical.

8. 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.

11. 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.
degrees between each hybrid orbital (like
a tetrahedron).

12. Hybrid orbitals can also contain non-bonding electrons
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

13. It is important to realize that sp3 hybridization is
just one possibility. Our course will include 5
major hybridization types:
Atomic Orbitals Used Hybrid Orbitals Formed
one s, one p two sp
one s, two p three sp2
one s, three p four sp3
one s, three p, one d five sp3d
one s, three p, two d six sp3d2

14. 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.

15. 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.

16. 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.

17. 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)

18. Lewis structures fail to explain how things like PCl5 are possible