I II III I. Lewis Diagrams (p. 184-189) Ch. 6 – Molecular Structure
Jan 21, 2016
I II III
I. Lewis Diagrams
(p. 184-189)
Ch. 6 – Molecular Structure
A. Octet Rule
Remember…· Most atoms form bonds in order to
have 8 valence electrons.
· Hydrogen 2 valence e-
· Groups 1,2,3 get 2,4,6 valence e-
· Expanded octet more than 8
valence e- (e.g. S, P, Xe)
· Radicals odd # of valence e-
Exceptions:
A. Octet Rule
F B FFH O HN O
Very unstable!!
F FF S FF F
B. Drawing Lewis Diagrams Find total # of valence e-. Arrange atoms - singular atom is
usually in the middle. Form bonds between atoms (2 e-). Distribute remaining e- to give each
atom an octet (recall exceptions). If there aren’t enough e- to go around,
form double or triple bonds.
B. Drawing Lewis Diagrams CF4
1 C × 4e- = 4e-
4 F × 7e- = 28e-
32e- FF C F
F
- 8e-
24e-
B. Drawing Lewis Diagrams BeCl2
1 Be × 2e- = 2e-
2 Cl × 7e- = 14e-
16e-
Cl Be Cl - 4e-
12e-
B. Drawing Lewis Diagrams CO2
1 C × 4e- = 4e-
2 O × 6e- = 12e-
16e-
O C O - 4e-
12e-
C. Polyatomic Ions
To find total # of valence e-:· Add 1e- for each negative charge.· Subtract 1e- for each positive
charge. Place brackets around the ion and
label the charge.
C. Polyatomic Ions
ClO4-
1 Cl × 7e- = 7e-
4 O × 6e- = 24e-
31e- OO Cl O
O
+ 1e-
32e-
- 8e-
24e-
NH4+
1 N × 5e- = 5e-
4 H × 1e- = 4e-
9e- HH N H
H
- 1e-
8e-
- 8e-
0e-
C. Polyatomic Ions
D. Resonance Structures
Molecules that can’t be correctly represented by a single Lewis diagram.
Actual structure is an average of all the possibilities.
Show possible structures separated by a double-headed arrow.
D. Resonance Structures
OO S O
OO S O
OO S O
SO3
I II III
II. Molecular Geometry
(p. 197-200)
Ch. 6 – Molecular Structure
A. VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
Electron pairs orient themselves in order to minimize repulsive forces.
A. VSEPR Theory
Types of e- Pairs· Bonding pairs - form bonds· Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
A. VSEPR Theory Lone pairs reduce the bond angle
between atoms.
Bond Angle
Draw the Lewis Diagram. Tally up e- pairs on central atom.
· double/triple bonds = ONE pair Shape is determined by the # of
bonding pairs and lone pairs.
Know the 8 common shapes & their bond angles!
B. Determining Molecular Shape
C. Common Molecular Shapes
2 total
2 bond
0 lone
LINEAR180°BeH2
3 total
3 bond
0 lone
TRIGONAL PLANAR
120°
BF3
C. Common Molecular Shapes
C. Common Molecular Shapes
3 total
2 bond
1 lone
BENT
<120°
SO2
4 total
4 bond
0 lone
TETRAHEDRAL
109.5°
CH4
C. Common Molecular Shapes
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
NH3
C. Common Molecular Shapes
4 total
2 bond
2 lone
BENT
104.5°
H2O
C. Common Molecular Shapes
PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
F P FF
D. Examples
CO2
O C O2 total
2 bond
0 loneLINEAR
180°
D. Examples
I II III
III. Polarity & IMF
(p. 204-207)
Ch. 6 – Molecular Structure
A. Dipole Moment
Direction of the polar bond in a molecule.
Arrow points toward the more e-neg atom.
H Cl+ -
B. Determining Molecular Polarity Depends on:
· dipole moments· molecular shape
B. Determining Molecular Polarity Nonpolar Molecules
· Dipole moments are symmetrical and cancel out.
BF3
F
F F
B
B. Determining Molecular Polarity Polar Molecules
· Dipole moments are asymmetrical and don’t cancel .
netdipolemoment
H2OH H
O
CHCl3
H
Cl ClCl
B. Determining Molecular Polarity Therefore, polar molecules have...
· asymmetrical shape (lone pairs) or · asymmetrical atoms
netdipolemoment
Dipole-Dipole Forces
Attractive forces between polar covalent molecules
London (Dispersion) Forces Attractive forces between the electron
clouds of large molecules in large quantity
Larger mass = Larger London Forces
Hydrogen Bonding
Special dipole-dipole attraction that involves H bonded with high electronegative elements N, O, or F