1. Section 6.5 – Molecular Geometry
The properties of molecules depend on the bonding and the molecular geometry, the 3- dimensional arrangement of the atoms in space.

2. Molecular Polarity
This is the uneven distribution of molecular charge, and it is determined by the polarity of each bond, along with the geometry of the molecule.

3. Two Theories – based on evidence
VSEPR Theory: Accounts for molecular bond angles.
Hybridization: Describes the orbitals that contain the valence electrons of a molecule’s atoms.

4. VSEPR Theory Valence Shell Electron Pair Repulsion
Repulsion between the valence-shell electrons surrounding an atom causes these sets to be oriented as far apart as possible.

5. CH4 NH3 H2O Draw the Lewis structures for all 3 molecules
Activity
CH4 NH3 H2O
Draw the Lewis structures for all 3 molecules
Sets of electrons that remain together in bonds or in lone pairs are referred to as electron domains. Electron domains prefer to be as far apart as possible from each other within a molecule.
How many electron domains are located around the central atom of each molecule?

6. Activity
CH4 NH3 H2O
Collect supplies – 1 set of gumdrops and toothpicks for each pair of students
Build a 3-D model of each molecule – SHOW ME YOUR MODEL!!!!
Remember – electron domains want to be as far apart as possible.
How do lone pairs affect the shape of each molecule?

7. Diatomic Molecules H2: non-polar HCl: polar
Diatomic molecules are composed of two atoms, so the geometry is always linear
the molecular polarity is determined by the electronegativity differences between the atoms.
H2: non-polar HCl: polar

8. Shorthand for Describing
For molecules containing more than two atoms, we can use the following symbols with subscripts:
A – the central atom
B – number of bonds
on the central atom
E – number of lone
pairs on central atom
(for atoms that have double or triple bonds, it is treated as a single B for geometry)

9. The Basis For VSEPR Theory
That one must consider the locations of all electron pairs of the valence electrons in the molecule.
Polyatomic ions are treated the same way.

10. The Basis For VSEPR Theory
The following examples do not have lone pairs that influence the geometry of the molecule.

11. Linear – AB2
Central atom with two single bonds, no lone pairs. Because the valence electron pairs in the bonds repel each other, the bonds are as far apart as possible (180°).
Ex: BeH2

12. Trigonal Planar – AB3
The 3 A-B bonds stay furthest apart by pointing to the corners of an equilateral triangle, giving 120° angles between the bonds.
Ex.: BH3

13. Tetrahedral – AB4
Octet rule is followed here. The distance between the A-B bonds is maximized if each bond points to the corners of a tetrahedron, giving bond angles of 109.5° between the bonds.
Ex.: CH4

14. Trigonal-bipyramidal – AB5
120° angles between bonds within the trigonal plane, 90° bond angles between the axial bond and those in the plane.
Ex.: PCl5

15. Octahedral – AB6
6 bonds to the central atom, all equidistant from each other. 90° bond angles.
Ex.: SF6

16. Lone Pairs Do Occupy Space and Influence Geometry
But our description of the molecular geometry refers to the positions of the atoms only.
A summary of the shapes of various molecules is in Table 6-5, p. 186.
Different sizes of B groups may distort some bond angles that are given in the table.

17. VSEPR and Unshared Electron Pairs
One must always write out the Lewis structure for a molecule to decide on the proper geometry, the chemical formula of something does not tell you about lone pairs around the central atom.

18. Bent – AB2E
2 bonds to central atom with one lone pair. The lone pair bends what one would expect to be linear. The lone pair takes up more space than a bond and shoves the bonded atoms closer together than the 120° for trigonal planar.

19. Bent – AB2E2
The addition of a second lone pair forces the bonding atoms even closer together than what one expects from tetrahedral.

20. Trigonal Pyramidal – AB3E
With 3 bonds one expects trigonal planar with 120° between the bonds, but the lone pair bends the plane away from the pair, forcing the atoms closer together.

21. Molecular Polarity Reflecting on the geometries, we can
now see why lone
pairs on the central
atom make a molecule polar – it changes the geometry of the molecule and creates an uneven “tug of war” across the molecule.

22. Intermolecular Forces
These are the forces of attraction that occur between molecules. They vary in strength but are generally weaker than regular bonding (ionic, covalent, or metallic).

23. Melting and Boiling Points
Usually are a good measure of the force of attraction. The higher the boiling point, the stronger the forces between particles.

24. Dipole Force The direction of the arrow
The strongest intermolecular forces exist between polar molecules. Each one acts as a dipole, created by equal but opposite charges that are separated by a short distance.
The direction of the arrow
is pointed to the negative pole,
the crossed tail indicates the + side

25. Dipole-dipole forces
These are the forces of attraction between polar molecules.
Example:
bp for F2 is -188°C
bp for HF is 20°C
bp for HCl is -85°C
Which is the stronger
dipole-dipole force?

26. For Molecules With More Than 2 Atoms
The molecular polarity depends on both the polarity and the orientation of each bond.

27. Induced Dipoles This can be very important in the
The electrons of a nonpolar molecule can be temporarily attracted by a polar molecule. This is weaker than a regular dipole-dipole force.
This can be very important in the
solubility of gases in water.

28. Usually Represented By
Hydrogen Bonding
A very special type of dipole-dipole force in which a hydrogen atom is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom of another molecule.
Usually Represented By
Dotted Lines

29. Explains High bp’s of Some Compounds
Gives the H atom a large positive charge, and it’s small size allows it to come very close to the unshared pair of electrons on an adjacent molecule.

30. Extremely Important in Biochemistry

31. Stereoisomers
Isomers – same chemical formula, different structures (isopropyl alcohol vs 2-propanol)
Stereoisomers – structures are mirror images.

32. WTF?????
Another World Trade Fund moment, brought to you by our sponsors, STERNO and Wally the Panda.

33. Think You Can Pass The Class?
Like a snowball’s chance in Hell!

34. London Dispersion Forces
This is a weak attractive force resulting from the imbalance of electrons and the creation of an instantaneous dipole. Important for noble gases and nonpolar molecules.

35. Increased Force
With the increased number of electrons in the interacting atoms or molecules, thus with increasing atomic or molar mass.

36. Assignment – Due Wed. EOP
Section 6.5 Worksheet
Mixed Review Worksheet
Molecular Geometry Worksheet
6.5 Textbook Problems