1. Chapter 9
Molecular Shape

2. VSEPR Theory Valence Shell Electron Pair Repulsion
Pairs of valence electrons are arranged as far apart from each other as possible.
Determines the shape of a molecule.

3. Hybrid Orbitals Orbitals of electrons in a bond
Combination of shapes and properties of the original atomic orbitals
sp, sp2, & sp3 hybridized

4. Bonding Orbital
To determine the number of bonding orbitals in a molecule, count the number of bonds around the CENTRAL ATOM of the molecule (things attached to the central atom)
Each bond will count for one bonding orbital
In this case multiple bonds will only count for one bonding orbital

5. Nonbonding Orbital
To determine the number of nonbonding orbitals in a molecule, count the number of lone electrons pairs around the CENTRAL ATOM of the molecule
Each PAIR of electrons will count for one nonbonding orbital

6. Steps for determining shape:
Draw the Lewis Dot Structure
Count the number of bonding orbitals
Count the number of nonbonding orbitals
Determine the shape
Redraw the Lewis Dot Structure to represent the shape

7. VSEPR Shapes We will learn about five shapes: Linear Trigonal Planar
Tetrahedral
Pyramidal
Bent

8. 1. Linear CO2 O = C = O 2 bonding orbitals 0 nonbonding orbitals
Bond angle = 180o
sp hybridized

9. 1. Linear

10. 2. Trigonal Planar BCl3 Cl B Cl Cl 3 bonding orbitals
0 nonbonding orbitals
Bond angle = 120o
sp2 hybridized

11. 2. Trigonal Planar

12. 3. Tetrahedral F CF4 F C F 4 bonding orbitals 0 nonbonding orbitals
Bond angle = 109.5o
sp3 hybridized

13. 3. Tetrahedral

14. 4. Pyramidal NH3 H N H H 3 bonding orbitals 1 nonbonding orbitals
Bond angle = 107o
sp3 hybridized

15. 4. Pyramidal

16. 5. Bent H2O H O H 2 bonding orbitals 2 nonbonding orbitals
Bond angle = 105o
sp3 hybridized

17. 5. Bent

18. Shape # orbitals b.o. nb. o. Bond <
linear 2
180o
trigonal planar 3
120o
bent 4
105o
pyramidal 1
107o
tetrahedral
109.5o

19. Lewis Structures in the Ghetto
Carbon makes 4 bonds, no unshared pairs
Nitrogen: 3 bonds, 1 unshared pair
Oxygen: 2 bonds, 2 unshared pairs
Hydrogen: 1 bond, no unshared
Halogens: 1 bond, 3 unshared pairs

20. Bond Length Different pairs of atoms form different bond lengths
As you go down a group, bond length increases
b/c atoms are getting larger
Multiple bonds are shorter than single bonds
Single > double > triple

21. Polar Covalent Bonds Polar Bond - atoms don’t share electrons equally.
Example: When O bonds with H, there is a difference in electronegativity between O & H which results in more of the electrons surrounding the O more frequently.
This causes the O end of the bond to be more negative and the H end of the bond to be more positive.

22. δ- O – H δ+ Polar Covalent Bonds
lower case Greek letter delta indicates charge
This bond contains a dipole, which merely means a positive and a negative end.

23. Electronegativity difference
In order to determine the polarity of a bond, calculate the difference in electronegativity
Electronegativity difference
Type of Bond
>2.0
Ionic
2.0 – 0.4
Polar Covalent
<0.4
Nonpolar

24. Polar molecules Molecules can be polar just like bonds
The polarity of a molecule is determined by:
The polarity of the individual bonds
The shape of the molecule

25. Polar molecules
If a molecule is symmetrical (in all 3 planes) it is nonpolar
If a molecule is asymmetrical (not symmetrical) it will be polar

26. Polar Mcl Exception
Tetrahedral Mcls are nonpolar when all four atoms attached to the central atom are the same.
CH4 – nonpolar
CH3Cl - polar

27. Polar Mcl Exception In the Ghetto
Tetrahedral, linear, and trigonal planar are usually nonpolar if the same thing is attached all the way around (like CH4 )
Pyramidal and bent are inherently asymmetrical and are usually polar

28. Shape Exceptions B & Al will form Trigonal Planar molecules
BF3 & AlH3
Be will form linear molecules
BeCl2