1. Molecular Shapes

2. Molecular Shapes
In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimized this repulsion.
We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

3. The VSEPR Model – common Mol. Geo.

4. Electron-Domain Geo.

5. The VSEPR Model
Predicting Molecular Geometries

6. The VSEPR Model
Predicting Molecular Geometries

7. The VSEPR Model
Predicting Molecular Geometries

8. The VSEPR Model
Molecules with Expanded Valence Shells

9. The VSEPR Model
Molecules with Expanded Valence Shells

10. The VSEPR Model
We determine the electron domain geometry by looking at electrons around the central atom.
We name the molecular geometry by the positions of atoms.
We ignore lone pairs in the molecular geometry.

11. The VSEPR Model
The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles
By experiment, the H-X-H bond angle decreases on moving from C to N to O:
Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.
Therefore, the bond angle decreases as the number of lone pairs increase.

12. The VSEPR Model
The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles
Similarly, electrons in multiple bonds repel more than electrons in single bonds.

13. The VSEPR Model Molecules with More than One Central Atom
In acetic acid, CH3COOH, there are three central atoms.
We assign the geometry about each central atom separately.

14. Polarity of Molecules Polar molecules interact with electric fields.
If the centers of negative and positive charge do not coincide, then the molecule is polar.

15. Polarity of Molecules Dipole Moments of Polyatomic Molecules
Example: in CO2, each C-O dipole is canceled because the molecule is linear. In H2O, the H-O dipoles do not cancel because the molecule is bent.

16. Polarity of Molecules
Dipole Moments of Polyatomic Molecules

17. Covalent Bonding and Orbital Overlap
Lewis structures and VSEPR do not explain why a bond forms.
How do we account for shape in terms of quantum mechanics?
What are the orbitals that are involved in bonding?
We use Valence Bond Theory:
Bonds form when orbitals on atoms overlap.
There are two electrons of opposite spin in the orbital overlap.

18. Covalent Bonding and Orbital Overlap

19. Covalent Bonding and Orbital Overlap

20. Hybrid Orbitals sp Hybrid Orbitals
Consider the BeF2 molecule (experimentally known to exist):
Be has a 1s22s2 electron configuration.
There is no unpaired electron available for bonding.
We conclude that the atomic orbitals are not adequate to describe orbitals in molecules.
We know that the F-Be-F bond angle is 180 (VSEPR theory).
We also know that one electron from Be is shared with each one of the unpaired electrons from F.

21. Formation of sp Hybrid Orbital
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Formation of sp Hybrid Orbital
A 2s orbital is mixed with one of the 2p orbitals to give two new orbitals with identical energies. Each resulting sp hydrbid orbital has one small and one much larger lobe. The two sp hybrid orbitals have the same shape with their large lobes pointing in opposite directions. The two large lobes of the sp hybrid orbitals shown in the figure on the right can each overlap with an orbital from another atom to form a covalent bond.

22. Formation of sp2 Orbitals
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Formation of sp2 Orbitals
The hybridization of one s and two p orbitals resulting in the formation of three sp2 orbitals is illustrated. Each sp2 orbital has one large and one small lobe. The sp2 orbitals all lie in the same plane and the large lobes point to the corners of an equilateral triangle. Each lobe can form a bond with another atom.

23. Formation of sp3 Orbitals
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Formation of sp3 Orbitals
One s and three p orbitals can hybridize to form four sp3 orbitals. The large lobe of each orbital points toward the vertex of a tetrahedron. Each large lobe can form a bond with another atom. The angle between each vertex of the tetrahedron is 109.5o.

24. FG09_017.JPG
Bonding in H2O
The O atom has 6 valence shell electrons. When it reacts with other molecules it undergoes sp3 hybridization, resulting in a tetrahedral arrangement of its orbitals. Two of its hybridized orbitals are used to form bonds to H and the other two are nonbonding orbitals, resulting in a bent geometry of the H2O molecule.

25. Hybridization in Ethylene
FG09_021.JPG
The sp2 hybridization of the two C atoms in ethylene is illustrated. Two of the sp2 hybrid orbitals on each C atom overlap with the 1s orbitals of a H atom to form sigma bonds. The remaining sp2 hybrid orbital forms a sigma bond between the two C atoms. The unhybridized p orbitals are perpendicular to the internuclear axis and can overlap to form a pi bond.

26. TB09_005.JPG
Table 9.4 p 366
This table gives the atomic orbital set, the hybrid orbital set and its geometry, and examples for sp, sp2, sp3, sp3d and sp3d2 hybridized orbitals.

27. Hybrid Orbitals Summary To assign hybridization:
draw a Lewis structure;
assign the electron pair geometry using VSEPR theory;
from the electron pair geometry, determine the hybridization; and
name the molecular geometry by the positions of the atoms.

28. Pi Bond Formation in Ethylene
FG09_022.JPG
The diagram on the left shows the sigma bonds in C2H4 and the two unhybridized p orbitals. To give each C atom eight valence electrons the p orbitals overlap to form a pi bond, resulting in the sharing of four electrons between the two atoms. The figure on the right shows the two lobes of the pi bond located above and below the sigma bond.

29. Triple Bond in Acetylene
FG09_023.JPG
The molecular geometry of C2H2 is linear, suggesting sp hybridization. After hybridization, each C atom has two unhybridized 2p orbitals. A sigma bond is formed between the two C atoms using sp hybrid orbitals. The p orbitals are perpendicular to the bond axis and perpendicular to each other, allowing the overlap of the p orbitals to form two pi bonds. Three pairs of electrons are shared by the C atoms, resulting in a triple bond.

30. Bonding in Benzene FG09_028.JPG
The 120o bond angles in benzene suggest sp2 hybridization. The hybridized orbitals on each carbon atom are used to form one sigma bond to hydrogen and sigma bonds to two adjacent carbon atoms. In addition, each carbon atom has one unhybridized 2p atomic orbital.

31. Orbitals of Benzene FG09_029.JPG
The overlap of 2p atomic orbitals on adjacent carbon atoms in benzene can be represented as forming a pi bond localized between these two C atoms. There are two equivalent ways to make localized pi bonds for benzene. A third representation allows the pi orbitals to spread out over the entire molecule, resulting in delocalization of the pi bonds. Delocalization gives a better representation of the properties of benzene.

32. Molecular Orbitals
Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. (E.g. why does O2 interact with a magnetic field?; Why are some molecules colored?)
For these molecules, we use Molecular Orbital (MO) Theory.
Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.
Molecular orbitals:
each contain a maximum of two electrons;
have definite energies;
can be visualized with contour diagrams;
are associated with an entire molecule.

33. Molecular Orbitals
The Hydrogen Molecule

34. MO Electron Configurations
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MO Electron Configurations
The MO diagrams are given for B2, C2, and N2 which all have large 2s-2p interactions and for O2, F2, and Ne2, which have small 2s-2p interactions.