1. 9-7 Polar Bonds vs. Polar Molecules (Section 12.3)
And you

2. Polar Bond = anytime there is a difference in electronegativity between atoms.
Polar Molecule = anytime the molecule can be divided by a plane, with one side that is δ+ and the other that is δ-, then the molecule is “polar”. In a sense, the bond dipoles must add like vectors to create a “net dipole”, an overall separation of charge. Symmetrical molecules, like BF3 can have bond dipoles, but they may either cancel or all point in one direction.
No net dipole = nonpolar

3. No net dipole = nonpolar

5. 9-8 Hybridization
When chemists measure the bond lengths and strengths of the 4 C – H bonds in CH4, the bonds are equivalent in both length and strength.
The question is: “How could four equivalent bonds form when the four valence electrons in the carbon atom are arranged with 2 in the 2s orbital, 1 in the first 2p orbital and 1 in another 2p orbital?” Like this...

6. Valence e-s 2p ___ ___ ___

7. To answer the above question and to describe the molecular orbital overlap (bonding) between atomic orbitals, chemists use “hybridization”.
Hybridization refers to the combination of atomic orbitals (s and p) to form “hybrid orbitals.”
 2p ___ ___ ___ SP3 ___ ___ ___ ___
2s ___ hybridized orbitals

8. For example, to form 4 equivalent orbitals, an s is combined with 3 p orbitals to form 4 equivalent sp3 hybrid orbitals. These four orbitals form a tetrahedral arrangement around the central carbon atom.

9. Example CH4

10. NEXT EXSMPLE... C2H4
If three equivalent bonds exist, as in C2H4, then an s orbital combines with 2 p orbitals to form 3 equivalent sp2 hybid orbitals. These 3 orbitals arrange themselves in a trigonal planar orientation. The unhybridized p orbital contains the remaining valence electron. As we’ll see later, this is involved in creating double (and or triple) bonds

11. 2p ___ ___ ___ 2s ___ sp2 ___ ___ ___ p ___

12. Example C2H2
When two, equivalent bonds are required, an s orbital hybridizes with a p orbital to form 2 sp hybrid orbitals. The other two remaining p orbitals are unhybridized. The two sp orbitals arrange in a linear geometry, while the p orbitals are perpendicular to each other. This type of bond arrangement is a triple bond, as seen in C2H2.

13. p ___ ___ ___ sp ___ ___ p ___ ___ s ___
The overlap of orbitals along the axis between the nuclei are called sigma (σ) bonds, whereas the bonds formed by the sideways overlap of p orbitals are termed pi (π) bonds.

14. One last thing here that relates hybridization to geometry:
sp hybridized atoms will have a linear geometry (2 e- density regions)
sp2 will have triangular plane geometry (3 e- density regions)
sp3 will have tetrahedral geometry (4 regions)
dsp3 will have triangular bi-pyramidal geo.(5 regions)
d2sp3 will have octahedral geo. (6 regions)
Notice that the d block electrons are incorporated into the structure, allowing the violation of the octet rule.
know this: single bond = σ bond
Double bond = σ and π bond
Triple bond = σ and 2π bonds