Chemical
Bonding
Mr. Matthew Totaro
Legacy High School
Honors Chemistry
"Perhaps one of you gentlemen would mind telling me just
what it is outside the window that you find so attractive...?"
2
Bonding Theories
• Bonding is the way atoms attach to make molecules.
• A chemical bond occurs when valence electrons are
either transferred or shared between the nuclei of
two atoms.
3
Chemical Bonds
Sodium to Chlorine Bond
4
Types of Bonds
5
Types of Bonds
• We can classify bonds based on the kinds of
atoms that are bonded together
Types of Atoms
metals to
nonmetals
nonmetals to
nonmetals
metals to
metals
Type of Bond
Ionic
Covalent
Metallic
Bond
Characteristic
electrons
transferred
electrons
shared
electrons
pooled
6
Bond Polarity
• The larger the difference in how strong one
atom pulls on the valence electrons, the more
polar the bond is.
Negative end toward more stronger atom.
d+ H — F d-
7
Bond Polarity
• Bonding between unlike atoms results in unequal
sharing of the electrons.
One atom pulls the electrons in the bond closer to its
side.
One end of the bond has larger electron density than
the other.
• The result is bond polarity.
The end with the larger electron density gets a partial
negative charge and the end that is electron deficient
gets a partial positive charge.
d+ H •• Cl d-
8
Types of Covalent Bonds
• Nonpolar (Pure) Covalent Bond
e- are shared equally
symmetrical e- density
usually identical atoms
Types of Covalent Bonds
• Polar Covalent Bond
e- are shared unequally
asymmetrical e- density
results in partial charges (dipole)
+
d
d
Bond Polarity
Nonpolar
Polar
Ionic
.
Electronegativity
• Measure of the pull an atom
has on bonding electrons.
• Increases across the period
(left to right).
• Decreases down the group
(top to bottom).
• The larger the difference in
electronegativity, the more
polar the bond.
Negative end toward more
electronegative atom.
12
Electronegativity, Continued
2.1
1.0 1.5
2.0 2.5 3.0 3.5 4.0
0.9 1.2
1.5 1.8 2.1 2.5 3.0
0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8
0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5
0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2
0.7 0.9 1.1
13
Electronegativity, Continued
14
Electronegativity and Bond
Polarity
• If the difference in electronegativity between
bonded atoms is 0 to 0.49, the bond is pure
(nonpolar) covalent.
• If the difference in electronegativity between
bonded atoms 0.5 to 1.7, the bond is polar
covalent.
• If the difference in electronegativity between
bonded atoms larger than 1.7, the bond is ionic.
15
Bond Polarity
3.0-3.0
= 0.0
3.0-2.1
= 0.9
Covalent
Nonpolar
0
0.49
3.0-1.0
= 2.0
Ionic
Polar
1.7
Electronegativity difference
4.0
16
Dipole Moments
• A dipole is a material with positively and
negatively charged ends.
• Polar bonds or molecules have one end slightly
positive, d+, and the other slightly negative, d-.
Not “full” charges, come from nonsymmetrical
electron distribution.
• Dipole moment, m, is a measure of the size of
the polarity.
Measured in debyes, D.
17
For Each of the Following Bonds, Determine
Whether the Bond Is Ionic or Covalent. If
Covalent, Determine if It Is Polar or Pure. If Polar,
Indicate the Direction of the Dipole.
• Pb-O
• P-S
• Mg-Cl
• H-O
18
For Each of the Following Bonds, Determine
Whether the Bond Is Ionic or Covalent. If
Covalent, Determine if It Is Polar or Pure. If Polar,
Indicate the Direction of the Dipole, Continued.
• Pb-O (3.5 - 1.9) = 1.6 \ polar covalent.
• P-S (2.5 - 2.1) = 0.4 \ nonpolar covalent.
• Mg-Cl (3.0 - 1.2) = 1.8 \ ionic.
• H-O (3.5 - 2.1) = 1.4 \ polar covalent.
19
Lewis Theory
• Lewis bonding theory
emphasizes the importance
of valence electrons.
• Uses dots to represent
valence electrons either on
or shared by atoms.
• Arranges bonding between
atoms to attain certain sets
of stable valence electron
arrangements.
G.N. Lewis
(1875-1946)
20
Lewis Symbols of Atoms
• Also known as electron dot symbols.
• Uses symbol of element to represent nucleus
and inner electrons.
• Uses dots around the symbol to represent
valence electrons.
 Put 2 electrons on one side, then one on the other 3.
• Remember that elements in the same group have
the same number of valence electrons; therefore,
their Lewis dot symbols will look alike.
21
Lewis Symbols of Ions
• Cations have Lewis symbols without
valence electrons.
Lost in the cation formation.
They now have a full “outer” shell that was the
previous second highest energy energy level.
• Anions have Lewis symbols with 8 valence
electrons.
Electrons gained in the formation of the anion.
Li•
Li+
••
:F:
•
•• −
[:F:]
••
22
Practice—Write the Lewis Symbol
for Arsenic.
23
9.1
-atoms will tend to gain, lose, or share
electrons until their outer energy level contains
eight electrons (Noble Gas e- configuration), or
two like Helium
-this produces the maximum stability in an atom
+
sodium metal
chlorine gas
table salt
Lewis Bonding Theory
• Atoms bond because it results in a more stable
electron configuration.
• Atoms bond together by either transferring or
sharing electrons.
• Usually this results in all atoms obtaining an
outer energy level 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.
 H, Li & Be try to achieve the He electron arrangement.
27
Ionic Bonds
• Metal to nonmetal.
• Metal loses electrons to form cation.
• Nonmetal gains electrons to form anion.
• Ionic bond results from + to − attraction.
• Lewis theory allows us to predict the
correct formulas of ionic compounds.
28
Example—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
29
Practice—Use Lewis Symbols to Predict the Formula
of an Ionic Compound Made from Reacting a Metal,
Mg, that Has 2 Valence Electrons with a Nonmetal,
N, that Has 5 Valence Electrons.
30
Practice—Use Lewis Symbols to Predict the Formula of
an Ionic Compound Made from Reacting a Metal, M, that
Has 2 Valence Electrons with a Nonmetal, X, that Has 5
Valence Electrons, Continued.

M 

 X 

 M


M 

 X 

Mg3N2
31
Covalent Bonds
• Often found between two nonmetals.
• Atoms bonded together to form molecules.
 Strong attraction.
• Atoms share pairs of electrons to attain octets (or duets).
• Molecules generally weakly attracted to each other.
 Observed physical properties of molecular substance due to these
attractions.
32
Single Covalent Bonds
• Two atoms share one pair of electrons.
 2 electrons.
• One atom may have more than one single bond.
••
••
••
F
H•
H
•H
O H
••
F
••
••
F
••
••
••
••
• F
••
••
••
•O
••
••
•
F •
••
••
••
••
F
33
Double Covalent Bond
• Two atoms sharing two pairs of electrons.
4 electrons.
• Shorter and stronger than single bond.
••
•O
••
•
•
••
•O
••
O ••
•• O
O
O
34
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
35
Bonding and 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.
Bonding pairs
••
••
••
•• O ••
••
•• S •• O
••
Lone pairs
36
Lewis Structures for
Covalent Compounds
37
Lewis Structures
Lewis structures are representations of
molecules showing all electrons, bonding
and nonbonding.
Writing Lewis Structures
PCl3
1. Find the sum of valence
electrons of all atoms in
the polyatomic ion or
molecule.
Keep track of the electrons:
5 + 3(7) = 26
 If it is an anion, add one
electron for each negative
charge.
 If it is a cation, subtract
one electron for each
positive charge.
Writing Lewis Structures
2. The central atom is the
least electronegative
element that isn’t
hydrogen. Connect the
outer atoms to it by
single bonds.
Keep track of the electrons:
26 − 6 = 20
Writing Lewis Structures
3. Fill the octets of the
outer atoms.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2
Writing Lewis Structures
4. Fill the octet of the
central atom.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0
Writing Lewis Structures
5. If you run out of electrons
before the central atom has
an octet…
…form multiple bonds until
it does.
Write the Lewis structure of CO2.
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.
45
Write the Lewis structure for NO3─
46
Example NO3─ , Continued
3. Attach atoms with pairs of
electrons and subtract from
the total.
N=5
O3 = 3∙6 = 18
(-) = 1
Total = 24 e-
O

O — N — O
Electrons
Start 24
Used 6
Left 18
Tro's Introductory Chemistry,
Chapter 10
47
Example NO3─ , Continued
3. Complete octets, outside-in.
 Keep going until all atoms
have an octet or you run out
of electrons.
:
:
O :



O — N — O

N=5
O3 = 3∙6 = 18
(-) = 1
Total = 24 e-
Electrons
Start 24
Used 6
Left 18

:

Electrons
Start 18
Used 18
Left 0
48
Example NO3─ , Continued
5. If central atom does not have
octet, bring in electron pairs
from outside atoms to share.
 Follow common bonding patterns
if possible.
:
:

O :



O — N — O


:
:
:

O

|
O — N
:

O:

49
Exceptions to the Octet Rule
• H and 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 two electrons to form two single bonds.
• B shares three 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:
50
Odd Number of Electrons
Though relatively rare and usually quite
unstable and reactive, there are ions and
molecules with an odd number of electrons.
Fewer Than Eight Electrons
• Consider BF3:
 Giving boron a filled octet places a negative charge
on the boron and a positive charge on fluorine.
 This would not be an accurate picture of the
distribution of electrons in BF3.
Fewer Than Eight Electrons
Therefore, structures that put a double bond
between boron and fluorine are much less
important than the one that leaves boron with
only 6 valence electrons.
Fewer Than Eight Electrons
The lesson is: If filling the octet of the central
atom results in a negative charge on the central
atom and a positive charge on the more
electronegative outer atom, don’t fill the octet of
the central atom.
More Than Eight Electrons
• The only way PCl5 can exist is if phosphorus
has 10 electrons around it.
• It is allowed to expand the octet of atoms on the
third row or below.
 Presumably d orbitals in these atoms participate in
bonding.
Resonance
Resonance
This is the Lewis
structure we would
draw for ozone, O3.
+
-
Resonance
• But this is at odds
with the true,
observed structure of
ozone, in which…
…both O—O bonds
are the same length.
…both outer oxygens
have a charge of -1/2.
Resonance
• One Lewis structure
cannot accurately
depict a molecule such
as ozone.
• We use multiple
structures, resonance
structures, to describe
the molecule.
Drawing Resonance Structures
1. Draw first Lewis structure that
maximizes octets.
2. Move electron pairs from
outside atoms to share with
central atoms.
3. If central atoms, 2nd row, only
move in electrons, you can
move out electron pairs from
multiple bonds.
··
··O ··
·· O
··
··
·· O
··
N
··
·· O ··
N
··
O ··
··
··
O
··
60
Practice—Draw Lewis Resonance
Structures for CO32(C is Central with O Attached)
61
Molecular Geometry
62
Molecular Shapes
• The shape of a
molecule plays an
important role in its
reactivity.
• By noting the number
of bonding and
nonbonding electron
pairs we can easily
predict the shape of the
molecule.
What Determines the Shape of a
Molecule?
• Simply put, electron pairs,
whether they be bonding
or nonbonding, repel each
other.
• By assuming the electron
pairs are placed as far as
possible from each other,
we can predict the shape
of the molecule.
Electron Domains
• This molecule has
four electron
domains.
• We can refer to the
electron pairs as electron
domains.
• In a double or triple bond,
all electrons shared
between those two atoms
are on the same side of the
central atom; therefore,
they count as one electron
domain.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains
is the one that
minimizes the
repulsions
among them.”
Electron-Domain
Geometries
These are the
electron-domain
geometries for two
through six electron
domains around a
central atom.
Electron-Domain Geometries
• All one must do is
count the number
of electron
domains in the
Lewis structure.
• The geometry will
be that which
corresponds to
that number of
electron domains.
Molecular Geometries
• The electron-domain geometry is often not the
shape of the molecule, however.
• The molecular geometry is that defined by the
positions of only the atoms in the molecules, not
the nonbonding pairs.
Molecular Geometries
Within each
electron domain,
then, there might
be more than
one molecular
geometry.
Linear Electron Domain
• In this domain, there is only one molecular
geometry: linear.
• NOTE: If there are only two atoms in the
molecule, the molecule will be linear no matter
what the electron domain is.
Trigonal Planar Electron Domain
• There are two molecular geometries:
 Trigonal planar, if all the electron domains are
bonding
 Bent, if one of the domains is a nonbonding pair.
Tetrahedral Electron Domain
• There are three molecular geometries:
 Tetrahedral, if all are bonding pairs
 Trigonal pyramidal if one is a nonbonding pair
 Bent if there are two nonbonding pairs
Trigonal Bipyramidal Electron Domain
• There are four
distinct molecular
geometries in this
domain:
 Trigonal bipyramidal
 Seesaw
 T-shaped
 Linear
Octahedral Electron Domain
• All positions are
equivalent in the
octahedral domain.
• There are three
molecular
geometries:
 Octahedral
 Square pyramidal
 Square planar
Practice—Predict the Shape Around the
Central Atom
PCl3
76
Practice—Predict the Shape Around the
Central Atom
SiH4
77
Practice—Predict the Shape Around the
Central Atom
SiF5-
78
Molecular Polarity
79
Polarity of Molecules
• In order for a molecule to be polar it must:
1. Have polar bonds.
 Electronegativity difference—theory.
 Bond dipole moments—measured.
2. Have an asymmetrical shape.
 Vector addition.
• Polarity effects the intermolecular forces of
attraction.
80
Molecule Polarity, Continued
The H—O bond is polar. Both sets of
bonding electrons are pulled toward the
O end of the molecule. The net result is
a polar molecule.
81
Molecule Polarity
The O—C bond is polar. The bonding
electrons are pulled equally toward both O
ends of the molecule. The net result is a
nonpolar molecule.
82
Example—Determining if a
Molecule Is Polar
N
H
H
H
83
Determine if NH3 is polar.
84
Example:
Determine if NH3 is polar.
• Write down the given quantity and its units.
Given:
NH3
85
Example:
Determine if NH3 is polar.
Information:
Given: NH3
• Write down the quantity to find and/or its units.
Find:
If polar
86
Example:
Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
• Design a solution map.
Formula of compound
Molecular polarity
Lewis structure
Bond polarity
and Molecular shape
87
Example:
Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
• Apply the solution map.
 Draw the Lewis structure.
Write skeletal structure.
H
N
H
H
88
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
Example:
Determine if NH3 is polar.
• Apply the solution map.
 Draw the Lewis structure.
Count valence electrons.
H
N=5
H=3∙1
Total NH3 = 8
N
H
H
89
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
Example:
Determine if NH3 is polar.
• Apply the solution map.
 Draw the Lewis structure.
Attach atoms.
H
N
H
H
N=5
H=3∙1
Total NH3 = 8
Start 8 eUsed 6 eLeft 2 e-
90
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
Example:
Determine if NH3 is polar.
• Apply the solution map.
 Draw the Lewis structure.
Complete octets.
H
∙∙
N
H
N=5
H=3∙1
Total NH3 = 8
H
Start 2 eUsed 2 eLeft 0 e-
91
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
Example:
Determine if NH3 is polar.
• Apply the solution map.
 Determine if bonds are polar.
H
∙∙
N
H
H
Electronegativity
N = 3.0
H = 2.1
3.0 – 2.1 = 0.9
\ polar covalent
92
Example:
Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
• Apply the solution map.
 Determine shape of molecule.
H
∙∙
N
H
4 areas of electrons
around N;
H
3 bonding areas
1 lone pair
N
H
H
H
Shape = trigonal pyramid
93
Example:
Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and
shape → molecule polarity
• Apply the solution map.
 Determine molecular polarity.
Bonds = polar
Shape = trigonal pyramid
N
H
H
H
Molecule = polar
94
Practice—Decide Whether Each of the Following
Molecules Is Polar
EN
O = 3.5
N = 3.0
Cl = 3.0
S = 2.5
••
•O
•
••
N
••
Cl ••
••
••
•O
•
••
••
•O•
• •
S
••
O ••
96
Practice—Decide Whether the Each of the
Following Molecules Is Polar, Continued
••
N
••
•O
•
••
Cl ••
••
Trigonal
bent
3.0
3.0
Cl
N
••
•O
•
••
3.5
Polar
S
S
••
O ••
Trigonal
planar
3.5 O
O
1. Polar bonds, N—O
2. Asymmetrical shape
••
•O•
• •
2.5
O3.5
1. Polar bonds, all S—O
2. Symmetrical shape
Nonpolar
3.5 O
97
Molecular Polarity Affects
Solubility in Water
• Polar molecules are
attracted to other polar
molecules.
• Since water is a polar
molecule, other polar
molecules dissolve well in
water.
And ionic compounds as well.
98
Properties of Substances
Ionic Bond
99
Properties of Substances
Metallic Bond
100
Properties of Substances
Covalent Bond
101
Molecular Solids
•Bond (Covalent)
Nonmetal + Nonmetal
•Example
Ice Water, H2O
•Melting Point
Low to moderate
•Conductivity
 Solid = no
 Liquid = no
 Aqueous = no
•Malleability
brittle
Water, H2O
102
Metallic Solids
• Bond (Metallic)
 Metal + Metal
• Example
 Copper, Cu
• Melting Point
 Low to High
• Conductivity
 Solid = yes
 Liquid = yes
 Aqueous = yes
• Malleability
 malleable
Copper metal, Cu
103
Ionic Solids
• Bond (Ionic)
 Metal + Nonmetal
• Example
 Table Salt, NaCl
• Melting Point
 High
• Conductivity
 Solid = no
 Liquid = yes
 Aqueous = yes
• Malleability
 brittle
Table Salt, NaCl
104
Covalent Network Solids
• Bond (Covalent)
 Nonmetal + Nonmetal
• Example
 Diamond, C
• Melting Point
 Very High
• Conductivity
 Solid = no
 Liquid = no
 Aqueous = no
• Malleability
 Brittle, very hard
105