1. 8.1 Molecular Geometry
Properties of Molecular Compounds depend upon 2 main things
Bonding
Molecular Geometry – arrangement of atoms in space
Used to determine Molecular Polarity
VSEPR Theory is used to predict molecular Geometry

2. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s ?

3. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e- ?

4. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e- ?

5. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e-
The Greatest Distance Apart.

6. VSEPR Theory (Ves-per Ther ee)
Valence-Shell Electron-Pair Repulsion Theory
Electron pairs surrounding a central atom repel each other as far apart from each other as possible
Bonded Pairs – electrons in bonds
Unshared Pairs – unattached electrons
Only those around the central atom. A B
Unshared Pair
Bonded Atom

7. VSEPR & Molecular Geometry
Type of Molecule
Bonded Atoms
Unshared Pairs
Molecular Shape
AB2 2
Linear
AB2E 1
Bent – 120
AB3 3
Triangular Planar
AB4 4
Tetrahedral
AB3E
Triangular Pyramidal
AB2E2
Bent
AB5 5
Triangular Bipyramidal
AB6 6
Octahedral

8. Practice
Draw the Lewis structure then predict the molecular geometry of the structure:
SO2
SF6
NI3
CF4
SiO2
IF4 1-
BCl3
Trigonal Planer, Angular
Octahedral, Octahedral
Tetrahedral, Trigonal Pyramid
Tetrahedral, Tetrahedral
Linear, Linear
Octahedral, Square Planar
Trigonal Planar, Trigonal Planar

9.   N
I I I F
F C F
F F
F S F
O = Si = O

10. Molecular Polarity  +  - Polar Bond – uneven sharing of electrons
Oxygen has a greater electronegativity than carbon C O C
 +
 -

11. Molecular Polarity  +  - Non-Polar Molecule
Opposite Polarities Cancel each other out.
Non-Polar Molecule

12. Molecular Polarity N H
Overall Polarity
• •
Polar molecule

13. General Rule for determining Polarity.
Non-Polar
Zero Unshared Pairs
If all the bonded atoms are of the same type.
Polar
1 or More Unshared Pairs
Unless linear shape

14. Identify the polarity of the following:
F F
F S F F
F C F
  N
I I I F
F C F Cl
    O H
O = Si = O

15. Hybridization
Carbon’s Electron Configuration
.    
1s 2s p
Each single electron is a bonding site.
Carbon only shows 2 bonding sites.
How is it possible for Carbon to bond to 4 Hydrogen atoms in Methane(CH4)?

16. Hybridization
Carbon’s New Electron Configuration
    .
1s 2s p
Moves one of the 2s electrons into the empty 2p orbital.
Creating four bonding sites.  
So we would expect 3 bonds to be identical and one to be different!

17. Hybridization
So we would expect 3 bonds to be identical and one to be different! C H s p
However, we find them all to be identical!

18. Hybridization
Hybridization – mixing 2 or more different orbitals, forming new identical orbitals.
For Carbon
1 s orbital + 3 p orbitals = 4 sp3 orbitals C H
sp3 s

19. 10.4

20. 10.4

21. Possible Combinations
1 s +1 p = 2 sp orbitals
1 s +2 p’s = 3 sp2 orbitals
1s +3 p’s = 4sp3 orbitals
1 s +3 p’s +1 d = 5 sp3d orbitals
1 s +3 p’s +2 d’s = 6 sp3d2 orbitals

22. Simplify To determine hybrid orbitals for a molecule.1 2
    O H 3 4
To determine hybrid orbitals for a molecule.
1st count the number of bonded atoms & unshared pairs
2nd use the following chart, where the number is the count from step 1.
3rd use a superscript for a letter that appears more than once. Example H2O = 4  sppp  sp3
Example BF3= 3  spp  sp2 1
F F B F 2
123456
spppdd 3

23. Practice
Determine the hybridization and Polarity of the central atom for the following molecules:
1) CO2
2) SO2
3) SF6
4) SiO2
1) O = C = O , (sp hybrid) on C
F F
3) F S F , (sp3d2)
4) O = Si = O , (sp hybrid) on Si

24. Molecular Bonding Molecular Orbitals 2 Types of Molecular Bonds
Resulting orbital formed by the overlap of 2 atomic orbitals.
2 Types of Molecular Bonds
Sigma Bond
Molecular orbital that is formed along the bonding axis.
Pi Bond
Molecular orbital that is formed above or below the bonding axis.

25. Sigma Bond

26. Sigma with Pi bond

27. Sigma with 2 Pi bond

28. Sigma bond (s) – electron density between the 2 atoms
Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms
Sigma bond (s) – electron density between the 2 atoms
10.5

29. 10.5

30. Sigma (s) and Pi Bonds (p)
1 sigma bond
Single 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
10.5

31. Sigma and Pi Bonds To determine type of bonds in a molecule:
All single bonds must form, sigma bonds.
Double bonds have 1 sigma and 1 pi bond.
Triple bonds have 1 sigma and 2 pi bonds.
In summary every bond has to have 1 sigma.
Pi bonds are only found in multiple bonds.

32. 8-2 Reviewsheet #13 Draw the Lewis Structure.
Instructions
#13
Draw the Lewis Structure.
Identify the orientation and shape.
Determine polarity of the molecule.
Identify the hybridization.
Determine the total sigma and pi bonds.

33. Intermolecular Forces
Forces of attraction between molecules.
Vary in strength
Weaker than Intramolecular Forces (Bonds)
Strength of IMF’s is related to the substance’s boiling points
IMF  BP 
3 Types of IMF’s
Dipole-Dipole
Hydrogen Bonding
London Dispersion

34. Dipole-Dipole
Attraction between polar molecules + -

35. Hydrogen Bonding Particularly Strong Dipole-Dipole
Attraction between Hydrogen atoms & an unshared electron pair on a strongly electronegative atom in another molecule
F, O, N
Holds DNA Together

36. London Dispersion Forces
Attractions between molecules that have an induced dipole or a momentary dipole
Due to the constant motion of molecules Induced dipole

37. Bond Type Substance Boiling Point (°C) Noble Gas He -269 Non-Polar H2
Boiling Points & Bond Types
Bond Type
Substance
Boiling Point (°C)
Noble Gas He
-269
Non-Polar H2
-253
Cl2
-34
Br2 59
Polar
NH3
-33
H2S
-61
H2O
100
Ionic
NaCl
1413
Metallic W
5660

38. Bond Length Distance between two bonded nuclei. 2 Common Trends
Different atoms bond at different lengths, depending upon the bond energies and size of the atoms.
2 Common Trends
Group trend between bonded diatomic molecules, increase in length. (top to bottom)
Double bonds are always shorter than single bonds between the same two atoms. Likewise triple bonds are shorter than both double and single bonds.

39. Practice Which of the following bonds would you think is the longest?
F - F
Cl - Cl
I - I
Answer: I - I

40. Practice #2
Which of the following bonds would have the shortest bond length? Why?
C – C
C = C
Answer :