Ch. 7 & 8- Chemical Bonding
Introduction to Bonding
IV
A. Vocabulary
 Chemical Bond

attractive force between atoms or ions that
binds them together as a unit

bonds form in order to…
 decrease
 increase
potential energy (PE)
stability
A. Vocabulary
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2
A. Vocabulary
COMPOUND
2 elements
Binary
Compound
NaCl
more than 2
elements
Ternary
Compound
NaNO3
A. Vocabulary
ION
1 atom
2 or more atoms
Monatomic
Ion
Polyatomic
Ion
+
Na
NO3
-
B. Types of Bonds
IONIC
COVALENT
Bond
Formation
e- are transferred from
metal to nonmetal
e- are shared between
two nonmetals
Type of
Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Other
Properties
odorous
C. Bond Polarity
 Most bonds are a
blend of ionic and
covalent
characteristics.
 Difference in
electronegativity
determines bond
type.
C. Bond Polarity
 Electronegativity



Attraction an atom has for a shared pair of
electrons.
higher e-neg atom  lower e-neg atom +
C. Bond Polarity
 Electronegativity Trend Increases up
and to the right.
C. Bond Polarity
 Nonpolar Covalent Bond



e- are shared equally
symmetrical e- density
usually identical atoms
C. Bond Polarity
 Polar Covalent Bond



+

e- are shared unequally
asymmetrical e- density
results in partial charges (dipole)

C. Bond Polarity
 Nonpolar
 Polar
 Ionic
View Bonding Animations.
C. Bond Polarity
Examples:
3.0-3.0=0.0
 Cl2
Nonpolar
 HCl
3.0-2.1=0.9
Polar
 NaCl
3.0-0.9=2.1
Ionic
Molecular Compounds
IV
A. Energy of Bond Formation
 Potential Energy


based on position of an object
low PE =
high stability
A. Energy of Bond Formation
 Potential Energy Diagram
attraction vs. repulsion
no interaction
increased
attraction
A. Energy of Bond Formation
 Potential Energy Diagram
attraction vs. repulsion
increased
repulsion
balanced attraction
& repulsion
A. Energy of Bond Formation
 Bond Energy

Energy required to break a bond
Bond
Energy
Bond
Length
A. Energy of Bond Formation
 Bond Energy

Short bond = high bond energy
B. Lewis Structures
 Electron Dot Diagrams




2s
show valence e- as dots
distribute dots like arrows
in an orbital diagram
4 sides = 1 s-orbital, 3 p-orbitals
EX: oxygen
2p
O
X
B. Lewis Structures
 Octet Rule


Most atoms form bonds in order to obtain 8
valence eFull energy level stability ~ Noble Gases
Ne
B. Lewis Structures
 Nonpolar Covalent - no charges
 Polar Covalent - partial charges
+
-

+
On Board Explanations
 Give examples of
 CaO
 O2F
 F2
C. Molecular Nomenclature
 The Seven Diatomic Elements
H
Br2 I2 N2 Cl2 H2 O2 F2
N O F
Cl
Br
I
IONIC Nomenclature
Monoatomic Ions – Single Element


Cations (loss of Electrons) - Metals
(positive)
Anions (gain of Electrons) - Non Metals
(negative)
Polyatomic Ions – Two or more Elements
Molecular Nomenclature
 Use Prefix
 Mono -1
 Di - 2
 Tri - 3
 Tetra - 4
 Penta - 5
Hexa
Hepta
Octa
Nona
Deca
-6
-7
-8
-9
- 10
Naming Molecular compounds
 1. Use prefix for both elements
 2. Last element ends in “ide”
 Exception to the rule
The prefix “Mono” is used for the second element
not the first
Ex: CO and CO2
Diatomic molecules names are NOT changed
II. Molecular
Geometry
IV
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
B. Determining Molecular Shape
 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!
C. Common Molecular Shapes
2 total
2 bond
0 lone
BeH2
LINEAR
180°
C. Common Molecular Shapes
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
C. Common Molecular Shapes
3 total
2 bond
1 lone
SO2
BENT
<120°
C. Common Molecular Shapes
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°
C. Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°
C. Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°
C. Common Molecular Shapes
5 total
5 bond
0 lone
PCl5
TRIGONAL BIPYRAMIDAL
120°/90°
C. Common Molecular Shapes
6 total
6 bond
0 lone
SF6
OCTAHEDRAL
90°
D. Examples
 PF3
4 total
3 bond
1 lone
F P F
F
TRIGONAL
PYRAMIDAL
107°
D. Examples
 CO2
2 total
2 bond
0 lone
O C O
LINEAR
180°
Rules for Lewis Dot structures
 Single Bond – Share 2 electrons
 Double Bond – Share 4 electrons
 Triple Bond – Share 6 electrons
 Coordinate Covalent Bond- Both Electrons
come from one single atom
 Octet Rule – All Atoms want 8 electrons

(EXCEPT Hydrogen 2 e and Boron 6 e)
Rules
 Count the total Valence Electrons
 Place the least electro-negative atom on
the inside & most electro-negative atom
on the outside
 Minus the electrons used for bonding
 Remainder of electron
A) distribute on the outer atom first
(octet rule)
 B) remainder put on central atom
 C) if short of electron – double or triple
bond

For ION formationNegative ions – Add an electron
Positive ions – Minus an electron
Resonance Structure – one or more valid
Lewis Dot Structure (ex SO3)
Free Radicals - Odd number of electrons
(rare)
Expanded Octet – any molecules with above
8 electron, can have up to 12 electrons
Above period 3 ( Because of D shell)
Ex : NF3
Ex: CS2
Ex:
CN-
I. Intermolecular
Forces
IV
A. Definition of IMF
 Attractive forces between molecules.
 Much weaker than chemical bonds within
molecules.
B. Types of IMF
B. Types of IMF
 London Dispersion Forces
View animation online.
B. Types of IMF
 Dipole-Dipole Forces
-
+
View animation online.
B. Types of IMF
 Hydrogen Bonding
C. Determining IMF
 NCl3

polar = dispersion, dipole-dipole
 CH4

nonpolar = dispersion
 HF

H-F bond = dispersion, dipole-dipole,
hydrogen bonding