1. Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Chapter 10

2. Valence Shell Electrons
Valence Shell Electron Pair Repulsion (VSEPR) Theory
Valence Shell Electrons
the outer shell electrons of an atom, the ones involved in bonding
for a given elements, # of valence electrons = Group Number
C: Group 4A, 4 valence electrons
O: Group 6A, 6 valence electrons
Draw Lewis structure
NH3

3. – Key: Electrons are all negatively charged (e-)
– Action: electron pairs (bonding pairs & lone pairs ) around the central
atom repel each other to keep themselves as far away as possible
– Outcome: Maximum separation
Minimum repulsion
Symmetry (equal repulsion)
– Applications: predict the electron pair geometry
predict the bond angle
predict the molecular geometry
predict the hybridization of the central atom
predict the polarity of the molecule
– Electron pair geometry : the arrangement of all electron pairs (bonding pairs & lone pairs ) surrounding the central atom of the molecule.
– Molecular geometry: the arrangement of only bonding pairs surrounding the central atom of the molecule

4. VSEPR CO2 Valence Shell Electron Pair Repulsion Example? O=C=O Class
AB2 2
Class
# of bonding
pairs
# of lone pairs
Electron pair
geometry
Molecular
Geometry
linear
Total # of
electron pairs
Example?
CO2
O=C=O

5. VSEPR # of bonding pairs # of lone pairs Total # of electron pairs
geometry
Molecular
Geometry
Class
trigonal planar
trigonal planar
AB3 3 3
trigonal planar
bent
AB2E 2 1 3

6. VSEPR # of bonding pairs # of lone pairs Total # of electron pairs
geometry
Molecular
Geometry
Class
tetrahedral
tetrahedral
AB4 4 4
tetrahedral
trigonal pyramidal
AB3E 3 1 4
tetrahedral
bent
AB2E2 2 2 4

7. # of bonding pairs # of lone pairs Total # of electron pairs
geometry
Molecular
Geometry
Class
trigonal
bipyramidal
trigonal
bipyramidal
AB5 5 5
trigonal
bipyramidal
seesaw
AB4E 4 1 5
trigonal
bipyramidal
T-shaped
AB3E2 3 2 5
trigonal
bipyramidal
linear
AB2E3 2 3 5

8. VSEPR # of bonding pairs # of lone pairs Total # of electron pairs
geometry
Molecular
Geometry
Class
octahedral
octahedral
AB6 6 6
octahedral
square pyramidal
AB5E 5 1 6
octahedral
square planar
AB4E2 4 2 6

9. Bond angle
Periodic table
# of valance shell electrons
Electron pair
geometry
Molecular
geometry
Lewis
structure
Total # of electron pairs
around central atom
Hybridization of
the central atom
What is the electron pair geometry, bond angle, molecular geometry and hybridization of the central atom of NH3?
electron pair geometry: tetrahedral
bond angle: 109.5o
molecular geometry: trigonal pyramidal
hybridization: sp3

10. bonding-pair vs. bonding
Repulsive force
bonding-pair vs. bonding
pair repulsion
lone-pair vs. lone pair
repulsion
lone-pair vs. bonding
>

11. Dipole Moments and Polar Molecules
Predicting Polarity of Molecules:
Dipole Moments and Polar Molecules
m = Q x r
Q is the charge
r is the distance between charges
1 D = 3.36 x C m H F
electron rich
region
electron poor d+ d-
• A measure of the polarity of a molecule
• The dipole moment can be determined experimentally
• Measure in Debye units
• Predict polarity by taking the vector sum of the bond dipoles

12. Molecules containing net dipole moments are called polar molecules
Molecules containing net dipole moments are called polar molecules. Otherwise they are called nonpolar molecules because they do not have net dipole moments.
Diatomic molecules: Determined by the polarity of bond
Polar molecules: HCl, CO,NO
Nonpolar molecules: H2,F2,O2
Molecules with three or more atoms: determined by the polarity of the bond and the molecular geometry
Dipole moment is a vector quantity, which has both magnitude and direction.

13. • Symmetric molecules such as these are nonpolar because the bond dipoles cancel
• All of the basic shapes are symmetric, or balanced, if all the domains and groups attached to them are identical
O C O C
O O
Linear molecule
Nodipolar moment
bent molecule
Net dipolar moment

14. Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4 O H S O
dipole moment
polar molecule
dipole moment
polar molecule C H C O
no dipole moment
nonpolar molecule
no dipole moment
nonpolar molecule

15. Molecular Geometry & Hybridization
• Lewis structures and VSEPR do not tell us why electrons group into domains as they do
• How atoms form covalent bonds in molecules requires an understanding of how orbitals interact
Molecular Geometry & Hybridization
Covalent Bonding Theories
• Valence bond (VB) theory
– Bonding is an overlap of atomic orbitals
• Includes overlap of hybrid atomic orbitals
• Molecular orbital (MO) theory
– Bonding happens when molecular orbitals are formed

16. Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals.
Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals.
Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.
Covalent bonds are formed by:
Overlap of hybrid orbitals with atomic orbitals
Overlap of hybrid orbitals with other hybrid orbitals

17. Predict the Hybridization of the Central Atom
Total # of
electron pairs
Electron pair
geometry
Hybridization
Examples 2
linear sp
BeCl2 3
trigonal planar
sp2
BF3 4
tetrahedral
sp3
CH4, NH3, H2O 5
trigonal bipyramidal
sp3d
PCl5 6
octahedral
sp3d2
SF6
Total # of electron pairs = number of hybrid orbital (sum of the superscripts!)

18. Formation of sp3 Hybrid Orbitals

19. Formation of sp Hybrid Orbitals

20. Hybridization in molecules containing double and triple bonds
Two Kinds of covalent Bonds
• Sigma bond(σ): bonding density is along the internuclear axis
– head to head overlap of two hybrid orbitals
– head to head overlap of p + hybrid
• any s character to the bond = sigma bond
• Pi bond(π): bonding density is above and below the
internuclear axis
– Sideway overlap p + p

21. (σ):
Sigma bond

22. (π):
Pi bond

23. Multiple bonds
• Double bond: One sigma and one pi
bond between the same atoms
–Example: ethylene
• Triple bond: One sigma and two pi
bonds between the same atoms
–Example: acetylene

24. Three Types of bonds:
1. Sigma bond(σ):
e-density (overlap region) is along the internuclear axis
2. Pi bond(π):
e-density (overlap region) is above and below the plane of
the internuclear axis.
Relative reactivity of bonds:
• Pi bonds are more reactive than sigma bonds
• Less energy is needed to break a pi bond than a sigma bond
3. Delocalized Bonds
• are not confined between two adjacent bonding atoms, but actually extend over three or more atoms.
• less reactive than normal pi bonds
• Example: benzene

25. Sigma (s) and Pi Bonds (p)
Framework of the molecule is determined by the arrangement of the sigma-bonds; Hybrid orbitals are used to form the sigma bonds and the lone pairs of electrons.
Sigma (s) and Pi Bonds (p)
Single bond
1 sigma bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? C H O
s bonds = 6
+ 1 = 7
p bonds = 1

26. Valence bond theory O
No unpaired e-
Should be diamagnetic
But experiment show that there are two unpaired electrons—paramagnetic.
Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals.

27. Molecular Orbital (MO) Theory
• Bonds are formed from interaction of atomic orbitals to form molecular orbitals.
MO theory takes the view that a molecule is similar to an atom
• The molecule has molecular orbitals that can be populated by electrons just like the atomic orbitals in atoms
• Number of MOs formed = number of atomic orbitals combined

28. Energy levels of bonding and antibonding molecular orbitals in hydrogen (H2).
A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed.
An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed.

29. Molecular Orbital (MO) Configurations
-MOs follow the same filling rules as atomic orbitals
The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined.
The more stable the bonding MO, the less stable the corresponding antibonding MO.
The filling of MOs proceeds from low to high energies.
Each MO can accommodate up to two electrons.
Use Hund’s rule when adding electrons to MOs of the same energy.(Electrons spread out as much as possible, with spins unpaired, over orbitals that have the same energy)
The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms.

30. ( ) - bond order = 1 2 Number of electrons in bonding MOs
Number of electrons in antibonding MOs ( - )
bond order 1