Chapter 6: Bonding…
Chemical
Bonding
 Describe covalent, ionic and metallic
bonds
 Classify bond type by electronegative
difference
 Explain why atoms form bonds
 Draw Lewis structures
 Classify covalent and ionic compounds
by physical properties
 Predict molecular geometry using VSEPR
theory
A. Vocabulary (Talkin’ the talk)
Chemical Bond
attractive force between atoms or ions
that binds them together as a unit
bonds form in order to…
decrease potential energy (PE)
increase stability
A. Vocabulary
CHEMICAL FORMULA
IONIC
Formula
Unit
NaCl
COVALENT
Molecular
Formula
(True molecules)
CO2
A. Vocabulary
ION – charged atom or
group of atoms
Positive
Charge
CATION
+2
Mg
NH4
+
Negative
Charge
ANION
-2
SO4
Cl
B. Basics of Ionic and Covalent Bonds
+
-
Na
Cl
IONIC
Electrons are: Transferred!
B. Ionic and Covalent traits
Ionic Bonding - Crystal Lattice
B. Ionic and Covalent traits
COVALENT
Cl
Cl
Electrons are: Shared!
B. Ionic and Covalent traits
Covalent Bonding - True Molecules
Diatomic
Molecule
B. Ionic and Covalent traits
C. Bond Polarity
OElectronegativity Trend (p. 151)
| Increases up and to the right.
C. Bond Polarity
 Electronegativity
1. Attraction an atom has for electrons.
2. In covalent molecules:
a) higher e-neg atom  b) lower e-neg atom +
C. Bond Polarity
Ionic
1.7
0.3
0.0
Polar
Covalent
Non polar
Covalent
% Ionic Character
4) Difference in
electronegativity
determines bond
type.
Difference in electronegativity
3) Most bonds are
a blend of ionic
and covalent
characteristics.
100%
3.3
50%
5%
0%
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
Elements
Electronegativity
difference
Bond type
More negative
atom
H&S
0.38
Polar
Cs & S
1.79
Ionic
S
S & Cl
0.58
Polar
Cl
C&H
0.35
Polar
C
O&O
0
Non-polar
none
H = 2.20
Cs = 0.79
S = 2.58
Cl = 3.16
S
C = 2.55 O = 3.44
D. Bond Formation
Potential Energy Diagram
attraction vs. repulsion
increased
repulsion
balanced attraction
& repulsion
D. Bond Formation
Potential Energy Diagram
attraction vs. repulsion
no interaction
increased
attraction
D. Bond Formation
Bond Energy
Energy required to break a bond
Bond
Energy
Bond
Length
E. Lewis Structures
Electron Dot Diagrams
show valence e- as dots
distribute dots like arrows
in an orbital diagram*
4 sides = 1 s-orbital, 3 p-orbitals*
Na
Mg
Al
C
N
O
F
X
Ne
E. Lewis Structures
Octet Rule
Most atoms form bonds in order to
obtain 8 valence eFull energy level stability ~ Noble
Gases
Ne
E. Lewis Structures
H H
|
|
|
|
H- C - C -H
H H
Structural
Formula
H H
H C C H
H H
Lewis Structure
What’s the difference?
E. Lewis Structures
 For covalent bonds
Element
Hydrogen
Group 2
Group 3 (B)
Carbon
Nitrogen
Oxygen
Halogens
Valence e-
1
2
3
4
5
6
7
Bonds Octet Rule
1
2
3
4
3
2
1
2
4 (rare)
6 (boron)
8
8
8
8
E. Lewis Structures
Examples
Covalent – show sharing of eIonic – show transfer of e-
+1 -1
Na
Cl → Na Cl
E. Lewis Structures
-1
Cl
Ca
Cl
+2 -1
→ Cl Ca Cl
E. Lewis Structures
Nonpolar Covalent - no charges
F
F → F F
Polar Covalent - partial charges
H
H
H
O
O →
δ+ H
δ-
E. Lewis Structures
Isomers – Same formula different
structure.
C3H7OH
F. Ionic and Covalent traits
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 mostly
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Make ‘formula units’
odorous, Make true
molecules
Other
Properties
G. VSEPR Theory
Valence Shell Electron Pair Repulsion
Theory
Electron pairs orient themselves in order
to minimize repulsive forces.
G. VSEPR Theory
Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
F. VSEPR Theory
 Use a chart to
determine shape
H. Common Shapes
2 Bonded to central atom
0 lone pairs
H Be H
BeH2
LINEAR
H. Common Shapes
3 Bonded to central atom
0 lone pairs
BF3
TRIGONAL PLANAR
F
F B F
H. Common Shapes
4 bonds to central atom
0 lone pair
H
H C H
H
CH4
TETRAHEDRAL
H. Common Shapes
3 bonded to central atom
1 lone pair
NH3
TRIGONAL PYRAMIDAL
H
H N H
H. Common Shapes
2 bonded to
central atom
2 lone
H O
H
H2O
BENT
I. Metallic Bonds
Metallic Bonding - “Electron Sea”
I. Metallic Bonds (on pg 5)
METALLIC
Bond
Formation
e- are delocalized
among metal atoms
Type of
Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water
no
Electrical
Conductivity
yes
(any form)
Other
Properties
malleable, ductile,
lustrous
I. Metallic Bonding
 Malleability
- Hammer into sheets
 Ductility
- Draw into thin wire
J. Intermolecular Forces
...that's all well and good, but what holds
different molecules together?
When H2O is in its solid state, what is
holding it like that?
When CO2 is in its solid state, what is
holding it like that?
J. Intermolecular Forces
Dipoles - polar covalent molecules where
there is an asymmetrical electron density
across the entire molecule (not just in the
bonds)
Dipole-dipole interaction occur when dipoles
line up and attract each other
Consider H2O, NH3, CH4, and CO2, given
the shape of the molecule, will one end
have a different charge than the other?
J. Intermolecular Forces
K. Hydrogen bonds
Hydrogen bonds - super dipole-dipole
interactions between δ+ on hydrogen and
δ- charges on other parts of a molecule
K. Hydrogen bonds - Ice
L. London Dispersion forces
What about non-polar molecules and atoms,
like the noble gases? How do they form
things like liquid helium and solid oxygen
without dipole interactions?
London dispersion forces are the attractive
forces between temporary dipoles created
by random motion of electrons
The more e- involved, the stronger the
London dispersion forces
L. London Dispersion forces of Helium