1. Molecular Structures Chapter 9
Linus Pauling ** Defined electronegativity.
Wrote “The Nature of the Chemical Bond” 1939.
Pioneered crystal and protein structures. 1 1 1 1

2. Molecular Shapes
Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which.
Lewis structures do not give us information about geometry or shape of molecules. H
H C H
For example, methane, CH4, is drawn as:
Where all the angles appear to be 90o. H C
Actually, the bond angles are 109.5o,
(tetrahedral angle)

3. Molecular Shapes Methane: Tetrahedral geometry Tetrahedron
(sp3-hybridization)

4. Some Molecular Geometries

5. 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.
This is the idea behind Valence Shell Electron Pair Repulsion (VSEPR) theory.

6. . . H : O : H Consider water, H2O Dot structure is:
TED=4 tells us that
electron domain geometry (EDG) is tetrahedral and,
BD: 2
NBD: 2
TED: 4
BD=2 tells us that:
molecular geometry (MG)
is bent

9. PCl5
SF6

10. Some examples: PBr3 SO3 CO32- SO32- H2O CH2=CH2 BeF2 BCl3 HCN
Work out the molecular geometries.

11. 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.

12. 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.

13. Polarity of Molecules
Dipole Moments of Polyatomic Molecules

14. Covalent Bonding and Orbital Overlap

15. Hybrid Orbitals

16. Hybrid Orbitals

17. Multiple Bonds
-Bonds: electron density lies on the axis between the nuclei.
All single bonds are -bonds.
-Bonds: electron density lies above and below the plane of the nuclei.
A double bond consists of one -bond and one -bond.
A triple bond has one -bond and two -bonds.

18. Multiple Bonds
Two p orbitals
overlap to form
a π-bond

19. Multiple Bonds Ethylene, C2H4, H2C=CH2, has:
one - and one -bond between the carbon atoms;
both C atoms sp2 hybridized;
both C atoms with trigonal planar molecular geometries.
2 p-orbitals
1 π-bond
5 σ-bonds

20. Multiple Bonds : N H C Consider acetylene, C2H2, H-CC-H, which has
one - and two -bonds between the carbon atoms;
both C atoms sp-hybridized;
both C atoms with linear molecular geometries.
When triple bonds form (e.g. HCCH, N2) one -bond is always above and below and the other is in front and behind the axis of the nuclei. N : C H
4 p-orbitals
2 π-bonds
3 σ-bonds
4 p-orbitals
2 π-bonds
1 σ-bond

21. Delocalized p Bonding In the case of benzene (C6H6) there are
6 C-C  bonds, all equal;
6 C-H  bonds, all equal.
Each C atom is sp2 hybridized.
There are 6 p orbitals on each C atom.
molecular
orbital
σ-bonds
p orbitals