1. Shapes

2. Writing Lewis Structures
Step 1
Total number of valence electrons Add the total negative charge if you have an anion. Subtract the charge if you have a cation.
Example: CO32- here it is
4 (for Carbon) + (3x6 oxygen's) + 2 charge = 24 total e-
  Step 2
Number of electrons if each atom is to be happy: Atoms in our example will need 8 e (octet rule) or 2 e ( hydrogen)
Step 3
Draw the structure: The central atom is C ( usually the atom with least electro negativity will be in the center). The oxygen's surround it . Because there are four bonds and only three atoms, there will be one double bond.
Step 4:
Octet the periphery (2 for hydrogen)
Left over for central atom = total (step 1) – Used (from step 3).
= 24 – (3x8) = 0
Step 5 
Double check your central atom has 8 electrons (except Boron), if not use multiple bonds for periphery to make it so .

3. Formal Charges: Writing Lewis Structures
Then assign formal charges.
For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
Subtract that from the number of valence electrons for that atom: the difference is its formal charge.
© 2009, Prentice-Hall, Inc.

4. Writing Lewis Structures
The best Lewis structure…
…is the one with the fewest charges.
…puts a negative charge on the most electronegative atom.
© 2009, Prentice-Hall, Inc.

5. Resonance
One Lewis structure cannot accurately depict a molecule like ozone.
We use multiple structures, resonance structures, to describe the molecule.

6. Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.

7. Molecular Shapes
The shape of a molecule plays an important role in its reactivity.
By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.
© 2009, Prentice-Hall, Inc.

8. What Determines the Shape of a Molecule?
Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.
© 2009, Prentice-Hall, Inc.

9. Electron Domains
We can refer to the electron pairs as electron domains.
In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.
The central atom in this molecule, A, has four electron domains.
© 2009, Prentice-Hall, Inc.

10. Valence Shell Electron Pair Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”
© 2009, Prentice-Hall, Inc.

11. Electron-Domain Geometries
These are the electron-domain geometries for two through six electron domains around a central atom.
© 2009, Prentice-Hall, Inc.

14. Hybridization Refers to mixing of orbitals.
Atomic orbitals of central atom undergo change to accommodate incoming atoms.
Hybridization could be sp, sp2, sp3, sp3d and sp3d2.
How do you tell the hybridization on the central atom?

15. 9.1 – 9.2: V.S.E.P.R. Valence-shell electron-pair repulsion theory
Because e- pairs repel, molecular shape adjusts so the valence e- pairs are as far apart as possible around the central atom.
Electron domains: areas of valence e- density around the central atom; result in different molecular shapes
Includes bonding e- pairs and nonbonding e- pairs
A single, double, or triple bond counts as one domain
Summary of LmABn (Tables ):
L = lone or non-bonding pairs
A = central atom
B = bonded atoms
Bond angles notation used here:
< xº means ~2-3º less than predicted
<< xº means ~4-6º less than predicted

16. Tables 9.1 - 9.3 # of e- domains & # and type of hybrid orbitals
e- domain geometry
Formula
Molecular geometry
Predicted bond angle(s)
Example (Lewis structure with molecular shape) 2
Two sp hybrid orbitals
Linear
AB2
180º
BeF2
CO2 |X X |B B A A

17. Three sp2 hybrid orbitals3
Three sp2 hybrid orbitals
Trigonal planar
AB3
120º
BF3
Cl-C-Cl
<< 120º
Cl2CO
LAB2
Bent
< 120º
NO21- |B B A |X X A |B : B A

18. Example: CH4 H | H—C—H Molecular shape = tetrahedral
Bond angle = 109.5º
109.5º

19. Four sp3 hybrid orbitals4
Four sp3 hybrid orbitals or
Tetrahedral
AB4
109.5º
CH4
LAB3
Trigonal pyramidal
< 109.5º
Ex: NH3 = 107º
NH3
L2AB2 Bent
<<109.5º
Ex: H2O = 104.5º
H2O B A A X X : A B X X A : A B

20. : PCl5 :Cl: :Cl: \ / :Cl—P—Cl: | :Cl:
\ /
:Cl—P—Cl: |
:Cl: :
Molecular shape = trigonal bipyramidal
Bond angles
equatorial = 120º
axial = 90º
90º
120º

21. Five sp3d hybrid orbitals5
Five sp3d hybrid orbitals
Trigonal bipyramidal
AB5
Equatorial = 120º
Axial = 90º
PCl5
LAB4
Seesaw
Equatorial < 120º
Axial < 90º
SF4 B A | A X | X X :
B - A - B B

22. Five sp3d hybrid orbitals5
Five sp3d hybrid orbitals
Trigonal bipyramidal
L2AB3
T-shaped
Axial << 90º
ClF3
L3AB2
Linear
Axial = 180º
XeF2 B : A | X A | : A B |

23. Six sp3d2 hybrid orbitals6
Six sp3d2 hybrid orbitals or
Octahedral
AB6
90º
SF6
LAB5
Square pyramidal
< 90º
BrF5 A X | A B | X X B B A B
| .. B B A X | X X

24. Six sp3d2 hybrid orbitals6
Six sp3d2 hybrid orbitals or
L2AB4
Square planar
90º
XeF4
L3AB3
T-shaped
<90º
KrCl31- A ..
| .. A B | .. B B B A ..
| .. A .. | B B .. B .. B