
 Molecular geometry : the three-dimensional arrangement of a molecule’s atoms in space

 “Valence-Shell Electron-Pair Repulsion”
 Refers to the repulsion between pairs of valence electrons of the atoms in a molecule
 VSEPR theory : repulsion between the sets of valencelevel electrons surrounding an atom causes these sets to be oriented as far apart as possible
 VSEPR theory is used to predict the shapes of molecules based on the fact that electron pairs strongly repel each other and tend to be oriented as far apart as possible


 VSEPR theory suggests that the lone pair occupies space around an atom just as bonding pairs do
 Actual shape of a molecule is determined by the positions of the atoms only
 Example: NH
3

 In VSEPR theory, double and triple bonds are treated in the same way as single bonds


 Atoms bonded to central atom: 2
 Lone pairs of electrons: 0
 Bond angles: 180°
 AX
2


 Atoms bonded to central atom: 3
 Lone pairs of electrons: 0
 Bond angles: 120°
 AX
3


 Atoms bonded to central atom: 2
 Lone pairs of electrons: 1
 Bond angles: <120°
 AX
2
E


 Atoms bonded to central atom: 4
 Lone pairs of electrons: 0
 Bond angles: 109.5°
 AX
4


 Atoms bonded to central atom: 3
 Lone pairs of electrons: 1
 Bond angles: 107°
 AX
3
E


 Atoms bonded to central atom: 2
 Lone pairs of electrons: 2
 Bond angles: 104.5°
 AX
2
E
2


 Atoms bonded to central atom: 5
 Lone pairs of electrons: 0
 Bond angles: 120°, 90°
 AX
5


 Atoms bonded to central atom: 6
 Lone pairs of electrons: 0
 Bond angles: 90°
 AX
6



Use VSEPR theory to predict the molecular geometry of AlCl
3
1. Draw the Lewis structure for the molecule.
.


2. Look at the number of bonds and lone electron pairs made in the molecule.
 How many bonds are attached to the central atom?
▪ Three
 How many lone electron pairs are there?
▪ None
3. Look at the type of molecule it is and refer to the chart to see what type of molecular shape it has.
 AlCl
3
= AX
3
= Trigonal-planar


 A) HI
▪ Linear
 B) CBr
4
▪ Tetrahedral
 C) AlBr
3
▪ Trigonal-planar
 D) CH
2
Cl
2
▪ Tetrahedral
 E) CO
2
▪ Linear
[double bond]





To explain how the orbitals of an atom can become arranged when the atoms form covalent bond, we use hybridization theory
 Hybridization : the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies
Hybrid orbitals : orbitals of equal energy produced by the combination of two or more orbitals on the same atom
The number of orbitals involved in hybridization determines the geometry of a molecule
There must be space around the atoms for every hybrid orbital
 Each orbital will be as far apart as possible from every other orbital

 Factors that affect the geometry of a molecule:
The number of bonds formed by each atom in the molecule
-
-
The number of lone pairs of electrons on the atoms
The sizes of the various types of atoms
The hybridization of some of the atoms’ orbitals

 Intermolecular forces : the forces of attraction between molecules
 Boiling point is a good measure of the force of attraction between particles of a liquid
 The higher the boiling point, the stronger the forces between particles


+ H Cl
 Molecular Polarity and Dipole-Dipole
Forces

-
 The strongest intermolecular forces exist between polar molecules
▪ Polar molecules act as tiny dipoles because of their uneven charge distribution


Dipole : created by equal but opposite charges that are separated by a short distance
Arrow points toward the more electronegative atom



Polar molecules contain permanent dipoles
Dipole-dipole forces : the forces of attraction between polar molecules
 These forces are short-range forces acting only between nearby molecules
 A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons

 Hydrogen Bonding:
 Hydrogen bonding : the intermolecular forces in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
 For a hydrogen bond to form, hydrogen must be bonded to either a fluorine, oxygen, or nitrogen atom
 The large electronegativity differences between hydrogen atoms and these atoms make the bonds connecting them highly polar

 London Dispersion Forces
 In any atom or molecule—polar or nonpolar– the electrons are in continuous motions
 At any instant the electron distribution may be slightly uneven
 creates a positive pole in one part of the atom and a negative pole in another
 London dispersion forces : the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles
 London forces act between all atoms and molecules
 They are only intermolecular forces acting among noble-gas atoms and nonpolar molecules

Directions: Determine the Lewis structure of the following molecules. Indicate the number of
atoms and number of lone pairs bonded to the central atom, and use the VSEPR theory to predict the molecular geometry shape of the respective molecules.




1. SeH
2
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
2. CF
4
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
3. PBr
3
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
4. BCl
3
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape




5. NI
3
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
6. CH
3
Br
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
7. HCN
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape
8. Na
2
O
▪ A) # bonded atoms
▪ B) # lone pairs
▪ C) Shape

Chemistry in Medicine
 Research the use of perfluorooctylbromide (C
8
F
17
Br) as an artificial oxygen carrier in synthetic blood. Construct an essay answering the following:
 What is perfluorooctylbromide?
 How does the compound work in the body to supply oxygen to tissues?
 What is the compound’s molecular structure?
 What types of bonds does the molecule contain?
 How many electron pairs do the C, F, & Br atoms in the compound share?