1. I. VSEPR = Valence Shell Electron-Pair Repulsion
Electron pairs repel each other
This applies to bonding pairs and lone pairs alike
Steps in Applying VSEPR
Draw the Lewis Structure
Count atoms and lone pairs and arrange them as far apart as possible
Determine the name of the geometry based only on where atoms are

2. Trigonal pyramidal Tetrahedral B. Complications to VSEPR
Lone pairs count for arranging electrons, but not for naming geometry
Example: NH3 (ammonia)
Lone pairs are larger than bonding pairs, resulting in adjusted geometries
Bond angles are “squeezed” to accommodate lone pairs
Trigonal pyramidal
Tetrahedral

3. b. Lone pairs must be as far from each other as possible
4. Double and triple bonds are treated as only 1 pair of electrons

4. linear trigonal planar bent tetrahedral bent trigonal bipyramidal.
5. Names for and examples of complicated VSEPR Geometries
linear
trigonal planar
bent
trigonal pyramidal
tetrahedral
bent
trigonal bipyramidal.
see-saw
T-shaped
linear
octahedral
square pyramidal
square planar

7. II. Hybridization Localized Electron Model sp3 Hybridization
Molecules = atoms bound together by sharing e- pairs between atomic orbitals
Lewis structures show the arrangement of electron pairs
VSEPR Theory predicts the geometry based on e- pair repulsion
Hybridization describes formation and properties of the orbitals involved in bonding more specifically
sp3 Hybridization
Methane, CH4, will be our example
Bonding only involves valence e-
a. H = 1s
b. C = 2s, 2p

8. If we use atomic orbitals directly, the predicted shape doesn’t match what we predict by VSEPR and what we observe experimentally
One C2s--H1s bond would be shorter than others
Three C2p--H1s bonds would be at right
angles to each other
We know CH4 is tetrahedral, symmetric
Free elements (C) use pure atomic orbitals
Elements involved in bonding will modify their orbitals to reach the minimum energy configuration

9. When C bonds to four other atoms, it hybridized its 2s and 2p atomic orbitals to form 4 new sp3 hybrid orbitals
The sp3 orbital shape is between 2s/2p; one large lobe dominates
When you hybridize atomic orbitals, you always get an equal number of new orbitals

10. The new sp3 orbitals are degenerate due to the mixing
The H atoms of methane can only use their 1s orbitals for bonding. The shared electron pair can be found in the overlap area of H1s—Csp3

11. NH3 and H2O also use sp3 hybridization, even though lone pairs occupy some of the hybrid orbitals.
C2H4, ethylene is our example
Lewis and VSEPR structures tell us what to expect
H atoms still can only use 1s orbitals
C atom hybridizes 2s and two 2p orbitals into 3 sp2 hybrid orbitals

12. The new sp2 orbitals are degenerate and in the same plane
One 2p orbital is left unchanged and is perpendicular to that plane
One C—C bond of ethylene forms by overlap of an sp2 orbital from each of the two sp2 hybridized C atoms. This is a sigma (s) bond because the overlap is along the internuclear axis.
The second C—C bond forms by overlap of the remaining single 2p orbital on each of the carbon atoms. This is a pi (p) bond because the overlap is perpendicular to the internuclear axis.

13. 9. Pictorial views of the orbitals in ethylene

14. sp Hybridization CO2, carbon dioxide is our example
Lewis and VSEPR predict linear structure
C atom uses 2s and one 2p orbital to make two sp hybrid orbitals that are 180 degrees apart
We get 2 degenerate sp orbitals and two unaltered 2p orbitals

15. Oxygen uses sp2 orbitals to overlap and form sigma bonds with C
Free p orbitals on the O and C atoms form pi bonds to complete bonding

16. d-orbitals can also be involved in hybridization
dsp3 hybridization in PCl5
d2sp3 hybridization in SF6

17. F. The Localized Electron Model
1. Draw the Lewis structure(s)
2. Determine the arrangement of electron pairs (VSEPR model).
3. Specify the necessary hybrid orbitals.

18. Polar Bonds
A. Any bond between atoms of different electronegativities is polar
Electrons concentrate on one side of the bond
One end of the molecule is (+) and one end is (-)
B. Complex molecules: vector addition of all bond dipoles

19. Polarity greatly effects molecular properties
Examples:
Polarity greatly effects molecular properties
Oppositely charged ends of polar molecules
Attract each other like opposite poles of magnets