Chapter 10
Chemical Bonding
Atoms interact with other atoms to form molecules, this is
chemical bonding
Bonding theories – are models that predict how atoms
bond together to form molecules
HIV-protease
protease inhibitor
Bonding Theories are
applied to design
molecules that will
interfere with the active
site of HIV-protease.
This delays or inhibits
the onset of AIDS.
CHAPTER OUTLINE








Chemical Bonds
Ionic Bonds and Covalent Bonds
Electronegativity
Bond Polarity & Electronegativity
Lewis Structures
Resonance
Molecular Shapes
Molecular Polarity
2
CHEMICAL BOND
 The
Mostnature
matterand
in nature
type of isthe
found
chemical
in form
bond
of is
directly
compounds:
responsible
2 or more
for elements
many physical
held together
and
chemical
through aproperties
chemical bond.
of a substance: (e.g. melting
conductivity).
 point,
Elements
combine together (bond) to fill their
outer energy levels and achieve a stable structure
(low energy).
 Noble gases are un-reactive since their energy
levels are complete.
3
CHEMICAL BOND
 This
When
difference
the conductivity
in
conductivity
apparatus is between
placed insalt
and
salt solution,
sugar is due
the to
bulb
the
different
will light.types of bonds
 between
But whentheir
it is atoms.
placed in
 Two
sugarcommon
solution,types
the bulb
of
bonding
does not are
light.
present:
ionic & covalent.
4
Lewis is known for:
Covalent bond
Lewis dot structures
Valence bond theory
Electronic theory of acids and
bases
Chemical thermodynamics
Heavy water
Named photon
Explained phosphorescence
Gilbert Newton Lewis (1875 - 1946) was a famous American physical chemist
known for the discovery of the covalent bond
(see his Lewis dot structures and his 1916 paper "The Atom and the Molecule")
Other major contributions were his theory of Lewis acids and bases and
Lewis coined the term "photon" for the smallest unit of radiant energy.
The Origin of Lewis
Symbols of Atoms
Drawings of cubical atoms, the
corners of the cube represented
possible electron positions
Lewis later cited these notes in
his classic 1916 paper on
chemical bonding, as being the
first expression of his ideas.
LEWIS
SYMBOLS OF ATOMS
 Lewis

In Lewis
structures
symbols
symbols,
foruse
valence
theLewis
firstelectrons
3symbols
periodsfor
toofshow
each
valence electrons
representative
element
are shown
elements
inas
molecules
a dot.
are shown
and ions
below:
of
compounds.
7
Lewis Bonding Theory
• atoms bond because it results in a more
stable electron configuration
• atoms bond together by either
transferring or sharing electrons so that
all atoms obtain an outer shell with 8
electrons
– Octet Rule
– there are some exceptions to this rule – the
key to remember is to try to get an electron
configuration like a noble gas
Lewis Symbols of Ions
• Cations have Lewis symbols without
valence electrons
– Lost in the cation formation
• Anions have Lewis symbols with 8
valence electrons
– Electrons gained in the formation of the
anion
Li•
Li+1
••
:F:
•
•• -1
[:F:]
••
Ionic Bonds
•
•
•
•
metal to nonmetal
metal loses electrons to form cation
nonmetal gains electrons to form anion
ionic bond results from + to - attraction
– larger charge = stronger attraction
– smaller ion = stronger attraction
• Lewis Theory allow us to predict the
correct formulas of ionic compounds
IONIC BOND
 After

Ionic bonds
bonding,
occur
each
between
when
atomelectrons
achieves
metals are
and
a complete
transferred
nonshell
metals.
between
(noble
twogas
atoms.
configuration).
Metal
Nonmetal
11
IONIC BOND
 Atoms

The smallest
that gain
lose
particles
electrons
electrons
of ionic
(metals)
(non-metals)
compounds
form form
positive
are
negative
ions
(cations).
(not
ions
atoms).
(anions).
Anion
Cation
12
Using Lewis Theory to Predict Chemical
Formulas of Ionic Compounds
Predict the formula of the compound that forms between
calcium and chlorine.
∙∙ ∙∙
∙
∙

Ca
  
: Cl : Ca2+
  


∙ Cl ∙∙
∙∙ ∙∙
∙ Cl ∙∙
∙∙ ∙∙
Transfer all the valance electrons
from the metal to the nonmetal,
adding more of each atom as you
go, until all electrons are lost
from the metal atoms and all
nonmetal atoms have 8 electrons
∙ Cl ∙∙
∙
∙
Ca
Draw the Lewis dot symbols
of the elements
  
: Cl :
  


CaCl2

Covalent Bonds
• typical of molecular species
• atoms bonded together to form
molecules
– strong attraction
• sharing pairs of electrons to attain
octets
• molecules generally weakly attracted to
each other
– observed physical properties of molecular
substance due to these attractions
COVALENT BOND
 Covalent
The smallest
bonds
particles
form when
of covalent
electrons
compounds
are
shared
are
molecules.
between two atoms.
 Covalent bonds form between two
Electrons
non-metals.
shared
15
Single Covalent Bonds
• two atoms share one pair of electrons
– 2 electrons
• one atom may have more than one single
bond
••
F
••
••
F
H•
H
•H
O H
••
F
••
••
F
••
••
••
••
• F
••
••
••
•O
••
••
•
F •
••
••
••
••
Double Covalent Bond
• two atoms sharing two pairs of electrons
– 4 electrons
• shorter and stronger than single bond
•
••
•O
••
•
••
•O
••
O •• O
••
O
O
Triple Covalent Bond
• two atoms sharing 3 pairs of electrons
– 6 electrons
• shorter and stronger than single or double
bond
••
•N
•
•
•
••
•N
•
N •• N
••
••
N
N
POLAR & NON-POLAR
BONDS
 Two types of covalent bonds exist:
Polar
&
Nonpolar
Electrons
shared
equally
 Non-polar covalent bonds occur between
similar atoms.
 In these bonds the electron
pair is shared equally
between the two protons.
19
POLAR & NON-POLAR
BONDS
 Polar covalent bonds occur between different
atoms.
 In these bonds the electron pair is shared
unequally between the two atoms.
 As a result there is a charge
separation in the molecule,
and partial charges on each
atom.
+ H F
20
Dipole Moments
• A dipole is a material with positively and
negatively charged ends
• Polar bonds or molecules have one end
slightly positive, +; and the other slightly
negative, – not “full” charges, come from nonsymmetrical
electron distribution
• Dipole Moment, m, is a measure of the size
of the polarity
– measured in Debyes, D
ELECTRONEGATIVITY
 Electronegativity
Linus Pauling derived
(E.N.)a is
relative
the ability of an
atom involved in aScale
covalent
bond
to attract
Electronegativity
based
on Bond
Energies.
the bonding electrons to itself.
Cs
F
0.7
4.0
Least
electronegative
Most
electronegative
22
ELECTRONEGATIVITY
23
BOND POLARITY &
ELECTRONEGATIVITY
Polarity is a measure of the inequality in the sharing of
bonding electrons
The more
different the
electronegativity
of the elements
forming the
bond
The larger the
electronegativity
difference
(EN)
The more
polar the
bond
formed
24
POLARITY &
ELECTRONEGATIVITY
As difference in
electronegativity
increases
Least
polar
Most
polar
Bond polarity
increases
25
POLARITY &
ELECTRONEGATIVITY
Electronegativity
difference
Bond Type
EN = 0
Non-polar covalent
0 < EN <1.7
Polar covalent
1.7 < EN
Ionic
26
POLARITY &
ELECTRONEGATIVITY
The molecule is
nonpolar covalent
Electronegativity
2.1
H
H
EN = 0
Electronegativity
2.1
Hydrogen Molecule
27
POLARITY &
ELECTRONEGATIVITY
The molecule is
polar covalent
+
H
Electronegativity
2.1
Cl
EN = 0.9
Electronegativity
3.0
Hydrogen Chloride Molecule
28
POLARITY &
ELECTRONEGATIVITY
No molecule exists
The bond is ionic
Na+
Electronegativity
0.9
ClEN = 2.1
Electronegativity
3.0
Sodium Chloride
29
SUMMARY
OF BONDING
EN > 1.7
Ionic Bond
(large EN)
Non-polar
(similar electronegativities)
EN = 0
Polar
(moderate EN)
Covalent Bond
(small to moderate EN)
0 < EN < 1.7
30
Bonding & Lone Pair Electrons
• Electrons that are shared by atoms are
called bonding pairs
• Electrons that are not shared by atoms
but belong to a particular atom are
called lone pairs
– also known as nonbonding pairs
LEWIS
STRUCTURES
 In a Lewis structure, a shared
electron pair is indicated by
two dots between the atoms, or
by a dash connecting them.
 Unshared pairs of valence
electrons (called lone pairs) are
shown as belonging to
individual atoms or ions.
32
LEWIS
STRUCTURES
 Structures
Covalent molecules
must satisfy
are octet
best represented
rule (8 electrons
with
around
electron-dot
each or
atom).
Lewis structures.
 Hydrogen is one of the few exceptions and forms
a doublet (2 electrons).
33
LEWIS
STRUCTURES
 Bonding
Non-bonding
electrons
electrons
can be
must
displayed
be displayed
by a dashed
as dots.
line.
34
Polyatomic Ions
• The polyatomic ions are attracted to
opposite ions by ionic bonds
– Form crystal lattices
• Atoms in the polyatomic ion are held
together by covalent bonds
Lewis Formulas of Molecules
• shows pattern of valence electron
distribution in the molecule
• useful for understanding the bonding in
many compounds
• allows us to predict shapes of molecules
• allows us to predict properties of
molecules and how they will interact
together
LEWIS
STRUCTURES
 More complex Lewis structures can be drawn by
following a stepwise method:
1. Count the number of electrons in the structure.
2. Draw a skeleton structure.
- most metallic element generally central
- halogens and hydrogen are generally
terminal
- many molecules tend to be symmetrical
- in oxyacids, the acid hydrogens are attached
to an oxygen
37
LEWIS
STRUCTURES
 More complex Lewis structures can be drawn by
following a stepwise method:
3. Connect atoms by bonds (dashes or dots).
4. Distribute electrons to achieve Octet rule.
5. Form multiple bonds if necessary.
38
Example 1:
Write Lewis structure for H2O
Step 1:
Step 2:
Step 3:
Step 4:
H2O
= 8 electrons
2 (1) + 6 = 8

H O H

Skeleton has
Hydrogen
structure
doublet
4 electrons used
Octet rule is satisfied should be
4 electrons remaining
symmetrical
39
Example 2:
Write Lewis structure for CO2
Step 1:
CO2
= 16 electrons 4 + 2(6) = 16
Step 2:
Step 3:
Step 4:
Step 5:






O C O 
Skeleton
structure
Octet
10
4 electrons
rule is NOT
used
Octet rule is satisfied should be
612electrons
electrons
satisfied
remaining
remaining
symmetrical
40
Writing Lewis Structures for
Polyatomic Ions
• the procedure is the same, the only
difference is in counting the valence
electrons
• for polyatomic cations, take away one
electron from the total for each positive
charge
• for polyatomic anions, add one electron to
the total for each negative charge
Example 3:
Write Lewis structure for CO32Step 1:
CO32-
= 24 electrons 4+3(6)+2 = 24
12
18
06 electrons
electrons
remaining



O

Step 5:

O C O

Step 4:


Step 3:

Step 2:
Octet
Octetrule
ruleisissatisfied
NOT
satisfied
42
Example 4:
Determine if each of the following Lewis structures
are correct or incorrect. If incorrect, rewrite the
correct structure.
Structure
has
14 electrons
Only
12
electrons
2(5) + 4(1) = 14
shown
2
4
2
2
2
Structure is
incorrect
Octets are complete
43
Exceptions to the Octet Rule
• H & Li, lose one electron to form cation
– Li now has electron configuration like He
– H can also share or gain one electron to have
configuration like He
• Be shares 2 electrons to form two single bonds
• B shares 3 electrons to form three single bonds
• expanded octets for elements in Period 3 or
below
– using empty valence d orbitals
• some molecules have odd numbers of electrons
– NO


:NO:
Some molecules, such as SF6 and PCl5 have more
than 8 electrons around a central atom in their
Lewis structure.
SF6 and PCl5 can violate the octet rule through the use of empty d orbitals:
both S and P can utilize empty d orbitals to hold pairs of electrons that help
bond halogen atoms.
Resonance
• we can often draw more than one valid
Lewis structure for a molecule or ion
• in other words, no one Lewis structure
can adequately describe the actual
structure of the molecule
• the actual molecule will have some
characteristics of all the valid Lewis
structures we can draw
Resonance
• Lewis structures often do not accurately represent
the electron distribution in a molecule
– Lewis structures imply that O3 has a single (147 pm) and
double (121 pm) bond, but actual bond length is
between, (128 pm)
• Real molecule is a hybrid of all possible Lewis
structures
• Resonance stabilizes the molecule
– maximum stabilization comes when resonance forms
contribute equally to the hybrid
O
O
+
O
O
O
+
O
Resonance
• we can often draw more than one valid
Lewis structure for a molecule or ion
• Real molecule is a hybrid
··
·
·
O
·
·
of all possible Lewis structures
·· O
··
represents resonance structures
··
·· O
··
The three oxygens are chemically equivalent, so it makes no
difference to the ion which oxygen assumes the double bond.
N
··
·· O ··
N
··
O ··
··
··
O
··
MOLECULAR
SHAPES
 The
A very
three-dimensional
simple model , VSEPR
shape of(Valence
the molecules
Shell
is an important
Electron
Pair Repulsion)
feature inTheory,
understanding
has beentheir
properties by
developed
andchemists
interactions.
to predict the shape of
molecules
basedhave
on their
Lewis
 large
All binary
molecules
a linear
shape since
structures.
they only contain two atoms.
 More complex molecules can have various
shapes (linear, bent, etc.) and need to be
predicted based on their Lewis structures.
49
MOLECULAR
SHAPES
 Based on VSEPR, the electron pair groups in a
molecule will repel one another and seek to
minimize their repulsion by arranging
themselves around the central atom as far apart
as possible.
 Electron pair groups can be defined as any one
of the following:
bonding pairs
non-bonding pairs
multiple bonds
50
SUMMARY OF
VSEPR SHAPES
Number of electron pair
groups around
central atom
Molecular
Shape
Bond
Angle
Examples
Bonding
Non-bonding
2
0
Linear
180
CO2
3
0
Trigonal planar
120
BF3
2
1
Bent
120
SO2
4
0
Tetrahedral
109.5
CH4
3
1
Pyramidal
109.5
NH3
2
2
Bent
109.5
H2 O
51
MOLECULAR
SHAPES
 Molecules with 2 electron pair groups around the
Linear molecules have polar bonds, but are
central atom form a linear shape.
usually non-polar.
Bond angle
is 180
Shape is
linear
2 electron pairs
around the
central atom
52
MOLECULAR
SHAPES
 Molecules with 3 electron pair groups around the
Trigonal planar molecules have polar bonds,
central atom form a trigonal planar shape.
but are usually non-polar.
Bond angle
is 120
3 electron
pairs
around the
central atom
Shape is
trigonal planar
53
MOLECULAR
SHAPES
 Molecules with 2 bonding pairs and 1 nonbonding
pairhave
groups
around
theand
central
atom
Bent
molecules
polar
bonds,
are polar.
form a bent shape.
1 Non-bonding
pair
Shape is bent
2 bonding pairs
around Bond
the angle
central atomis 120
54
MOLECULAR
SHAPES

Molecules
with 4 electron
pairsbonds,
groupsand
around
Tetrahedral
molecules
have polar
are
the central
atom form a tetrahedral shape.
usually
non-polar.
4 bonding pairs
around the
Shape is
central atom
tetrahedral
Bond angle
is 109.5
55
MOLECULAR
SHAPES
 Molecules with 3 bonding pairs and 1 nonPyramidal
molecules
havearound
polar bonds,
and are
polar.
bonding
pair groups
the central
atom
form a pyramidal shape.
Shape is
pyramidal
1 Nonbonding pair
3 bonding pairs
Bond angle
around
the
is 109.5
central
atom
56
MOLECULAR
SHAPES
 Molecules with 2 bonding pairs and 2 nonBent
molecules
polar
bonds,
are polar.
bonding
pairhave
groups
around
theand
central
atom
form a bent shape.
Shape is
bent
2 Non-bonding
pair
2 bonding pairs
Bond angle is
around the
109.5
central atom
57
SUMMARY OF
MOLECULAR SHAPES
Linear
Trigonal planar
Bent
Pyramidal
Tetrahedral
Symmetrical shapes
Polar bonds
Non-polar molecules
Unsymmetrical
shapes
Polar bonds
Polar molecules
58
Polarity of Molecules
• For a molecule to be polar it must
1) have polar bonds, symmetrical shape, and
different terminal atoms
2) have polar bonds
• electronegativity difference - theory
• bond dipole moments – measured
3) have an unsymmetrical shape
• using vector addition
• polarity effects the intermolecular forces
of attraction
Dipole moment is the measured polarity of a polar covalent bond.
It is defined as the magnitude of charge (electrons) on the atoms and the
distance between the two bonded atoms.

:OCO:
O
H

H
polar bonds,
and unsymmetrical
shape causes molecule
to be polar
polar bonds,
but nonpolar molecule
because pulls cancel
Cl
Cl
Cl
Cl
C
C
CCl4
m = 0.0 D
Cl
H
H
Cl
CH2Cl2
m = 2.0 D
Adding Dipole Moments
62
COMPARING PROPERTIES
OF IONIC & COVALENT
COMPOUNDS
Ionic
Covalent
Structural Unit
Ions
Molecules
Melting Point
High
Low
Boiling Point
High
Low
Solubility in H2O
High
Low or None
Electrical Cond.
High
None
NaCl, AgBr
H2, H2O
Examples
63
THE END
64