1. Hybridization

2. VSEPR Theory Review Valence electrons only are involved in bonding.
Non-bonding and bonding electron pairs around the central atom repel each other.
This repulsion causes specific shapes and bond angles for each molecule.

3. Question
How can atoms, that have s, p and d orbitals, bond in ways that make the molecule shapes?
Ex: p orbitals at 90o angles to each other.
How can they make a bond angle of 120o in trigonal planar?? ?

4. Energy
1s 2s p s
Carbon Hydrogen
How can carbon make four bonds with four hydrogen atoms?? How can those bonds be at 109.5o (tetrahedral)?

5. Hybridization: when hybrid orbitals form.
Answer: Hybrid orbitals: the sublevels in an atom’s outer shell recombine into new orbitals of equal energy with different shapes and angles.
Energy
1s sp3 hybrid orbital
Hybridization: when hybrid orbitals form.

6. Total energy the same, but redistributed
1s 2s p s
Carbon Hydrogen
Energy
1s sp3 hybrid orbital
Total energy the same, but redistributed 6

7. After hybridization Total energy is the same.
Energy redistributed equally among four new hybrid orbitals.
Hybrid orbitals are more directional: in CH4, they point out to the four corners of the tetrahedral shape.

8. Shapes of hybrid orbitals
s, p
s, p, p
trigonal pyramidal
s, p, p, p

9. NH3 is also sp3 hybridization: three orbitals for the H, one orbital for the non-bonding pair.

10. For each compound’s central atom, draw the orbital notation.
Then, draw the orbital notation after hybridization.
Name the type of hybridization and the shape of each molecule.
SiF4
CO2
BF3
PH3
H2O

11. Homework
P. 117 # 1 and 2

12. Multiple Bonds

13. When a double bond forms, the two bonds are not exactly the same.
First: “end on” s-orbital interaction: σ (sigma) bonds
Second: “side on” p-orbital interaction: π (pi) bonds

14. Single bonds: always σ bonds.
Double bonds: one σ + one π bond.
Triple bonds: one σ bond +two π bonds

15. Count sigma and pi bonds
CHCl3
SCO
SeO2
ClO3-

16. Practice Problems
Count the total number of sigma and pi bonds in each molecule
Draw a Lewis structure first!
H2CO O2
CO2
HCN (C is the central atom)
CSN- (C is the central atom)
N3-

17. Homework Draw Lewis structures and count sigma and pi bonds for: OCN-
NO2-
NO3- O3
SO3

18. Delocalization of Electrons

19. CO32- How do you figure out which oxygen to make the double bond??

20. You can draw the molecule THREE different WAYS!
Resonance Structures: Any of the Lewis structures that can be drawn if a double bond could be in more than one place.

21. Are any of these structures correct?
What would the bond lengths be like?
The double bond is shorter than the single bonds.
The bond lengths are not actually different!

22. But where are the electrons?
In a π bond (double or triple bonds only), the electrons can spread over more than 2 nuclei.

23. Why do π bonds delocalize?
When the electrons spread out, it gives the molecule a lower potential energy.
The molecule is more stable.

24. Remember: Resonance structures are imaginary.
The electrons are actually being shared by more than two atoms.
So we need more than one picture to show this.

25. Another Example: NO2-

27. If more than one Lewis structure can be drawn, what actually happens is part-way between those Lewis structures.
Resonance structures have the same σ bonds, but different π bonds from one another.

28. O3 Ozone

29. Which of the following has more than one possible Lewis structure?
NH4+
HCO3-
C2H2
OH-

30. Draw a Lewis structure for NO3-.
Does this really explain the shape of the ion? Why or why not?
Circle the electron pair that is delocalized.
Is it a sigma or a pi bond?
If there is one electron pair being shared by all the oxygen atoms, what is the charge of each oxygen atom? (Hint: the charge of the ion divided by the number of O atoms.)
Repeat 1-5 with CO32-