1. The Nature of the Chemical Covalent Bond
Section 4.2

2. Valence Bond Theory
A half-filled orbital in one atom can overlap with another half-filled orbital of a second atom to form a new bonding orbital
The new bonding orbital form the overlap of atomic orbitals contains a pair of electrons of opposite sign
The total number of electrons in the bonding orbital must be two
When atoms bond, they arrange themselves in space to achieve the maximum overlap of their half-filled orbitals. Maximum overlap produces a bonding orbital of lowest energy.

3. Hybrid Orbitals Hybridization:
A theoretical process involving the combination of atomic orbitals to create a new set of orbitals that take part in bonding
Hybrid orbitals only take place to form a new molecule, they do not exist on their own.
Two problems from Lewis theory: inability to explain the four equal bonds represented by the four pairs of electrons in a carbon cmpd & the existence of multiple bonds

4. Forms of Hybridization
Depending on the types of orbitals that overlap you will get different shapes of molecules:
sp hybridization occurs when an s and a p orbital overlap –two sp combine result = linear
sp2 hybridization occurs when an s and 2 p orbitals overlap, – three sp2 combine result = trigonal planar
sp3 hybridization occurs when one s and three p orbitals overlap – four sp3 combine result = tetrahedral

5. Single Covalent Bonds
Single bonds form from the overlap of end to end orbitals
The original bond formed between two atoms is called a sigma (denoted by the Greek lowercase s, ) bond

6. Double and Triple Covalent Bonds
Double and Triple bonds form from the side to side overlap of p orbitals over the sigma bond.
They are called pi (denoted by the Greek lowercase p, ) bonds, a double bond has one  bond and a triple bond has 2  bonds

7. Double bonds:
Triple Bonds:

8. Polar Covalent Bonds
Polar bonds happen when the differences in electronegativities between two non-metals is greater than 0.5

9. Polar Molecules
Occur when there are polar bonds AND there is an unequal sharing of the positive and negative sides.
If there are unpaired electrons, treat them as negative charges.

10. Polar or Not
CBr4
HBr
NF3
BF3
H2S