Chapter #10
Chemical Bonding
CHAPTER 12
Forces Between Particles
 Noble Gas Configurations
 Ionic Bonding
 Covalent Bonding
 VSEPR Theory and Molecular Geometry
 Electronegativity
 Polar Bonds and Molecules
Atomic Stability
It has been recognized for a long time that the noble gases have
great chemical stability. With few exceptions they are
unreactive or inert.
The noble gases have 8 valence electrons with the exception of
He which has 2.
He
Ne
Ar
Kr
Xe
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
Lewis Diagrams
The electronic configuration of the noble gases is described as
being energetically stable.
We can draw a Lewis diagram to illustrate the number of
valence electrons an atom has.
In a Lewis diagram valence electrons are represented by dots
placed above, below and to the left and right of the atoms
symbol.
e.g. element with 4 valence electrons
E
Lewis Diagrams
There are two simple rules to keep in mind when drawing
Lewis diagrams:
• Place one dot in each of the four locations before
doubling up.
• There can be only a maximum of 2 dots in any one
location.
E 
E 
E 
E 
Lewis Diagrams
What is the Lewis diagram for H?
1. First write the electron configuration:
1s1
2. Identify the number of valence electrons.
1 valence electron.
H
For a representative element it is easy to identify the number of
valence electrons as this is equal to the group number.
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
Lewis Diagrams
What is the Lewis diagram for S?
1. First write the electron configuration:
[Ne]3s23p4
2. Identify the number of valence electrons.
6 valence electrons
S
Alternatively you can recognize that S is in
group VIA so has six valence electrons
LEWIS STRUCTURES OF THE ELEMENTS
1
2
13
14
15
16
17
18
He
H
Be
B
C
N
O
F
Ne
Na Mg
Al
Si
P
S
Cl
Ar
Li
LEWIS STRUCTURES OF IONS
(AFTER REMOVAL OR ADDITION OF
ELECTRONS)
1
1+
2
H
2+
Li
Be
Na Mg
13
B
3+
Al
14
15
16
17
18
4-
3-
2-
1-
He
C
N
O
F
Ne
Si
P
S
Cl
Ar
Lewis Diagrams
The octet rule states that:
“Atoms interact in order to obtain a stable octet of eight
valence electrons”
The octet rule works extremely well at describing the
interactions of the representative elements.
Lewis Diagrams
One way in which atoms can interact to satisfy the octet rule is
by transferring electrons between each other.
Transferring of electrons results in the atoms acquiring net
positive and negative charges.
When an atom loses or gains electrons a simple ion is
formed.
Cations have more protons than electrons and are positive.
Anions have more electrons than protons and are negative.
Ion Formation
Consider a Na atom what happens if it loses one electron?
Na
I.E.
Na+
1e-
+
[Ne]
[Ne]3s1
11 P and 10 e-
11 P and 11 e-
Consider a Cl atom would you expect it to lose or gain
electrons?
Cl
[Ne]3s23p5
17 P and 17 e-
+
1e-
E.A.
Cl[Ne]3s23p6
17 P and 18 e-
Ion Formation
Metals tend to lose electrons forming positively charged ions
called cations.
• A representative metal will lose its group number of
electrons to obtain a stable octet.
Na
→
Na+
Mg
→
Mg2+ +
+
1e- ( Isoelectronic with Ne)
2e-
(isoelectronic with Ne)
What would the charge be of the ion formed by a Li atom?
And which Noble gas is it isoelectronic with?
+1 The ion formed would be Li+
Isoelectronic with He
Ion Formation
Non-metals tend to gain electrons forming negatively
charged ions called anions.
• A representative non-metal will gain (8 - group number)
electrons to obtain a stable octet.
O +
2e→
O2- (isoelectronic with Ne)
S +
2e→
S2- (isoelectronic with Ar)
What would the charge be of the ion formed by a I atom?
Which Noble gas is it isoelectronic with?
-1 The ion formed would be IIsoelectronic with Xe
Lewis Structure of NaCl
+
+
+
Na Cl Na Cl Na Cl
Cl- Na+Cl-Na+Cl-Na+
Forces between oppositely charged ions are called
Ionic bonds. Each ion is surrounded by an octet of
Electrons, thus making the ions stable.
Crystal Lattice of NaCl
Ionic compounds do not exist as discrete molecules. Instead
they exist as crystals where ions of opposite charges occupy
positions known as lattice sites.
Ions combine in the ratio
that results in zero
charge to form ionic
compounds.
Which ions are the
smaller ones?
Crystal Lattice of NaCl
Crystal Lattice of NaCl
Ionic compounds do not exist as discrete molecules. Instead
they exist as crystals where ions of opposite charges occupy
positions known as lattice sites.
Ions combine in the ratio
that results in zero
charge to form ionic
compounds.
Which ions are the
smaller ones? Sodium
Crystal Lattice of NaCl
Molecular Compounds
In our early lectures we defined a molecule as “as a
compound made of nonmetals.”
Molecules exist as particles containing the number of atoms
specified by their formula.
e.g. a water molecule is a particle containing 2 hydrogen
atoms and one oxygen atom and has the formula H2O.
Molecular Compounds
Non-metals may also complete their octets by sharing
electrons.
This may occur between non-metal atoms of the same type:
e.g. H2, O2, N2, Cl2, F2, I2, etc
Or between different types of non-metal atoms:
e.g. CO2, H2O, CH4, etc
Molecular Compounds
Consider two hydrogen atoms separated by a large
distance. Each has 1 electron in a 1s atomic orbital.
-
-
+
+
Why does the electron stay around the nucleus?
Now lets bring the two atoms together so there
orbitals overlap.
Molecular Compounds
-
+
+
-
The atomic orbitals overlap to form a new molecular
orbital. This is a stable configuration as each H atom
can have a full 1s subshell (like He) where the
electrons spend most of their time shared between
the atoms. In this arrangement each nucleus feels an
inwards attraction to the two electrons. This is called
covalent bonding.
Molecular Compounds
-
+
+
-
This new arrangement of protons and electrons is more
stable than separate hydrogen atoms since the
attraction of a proton to two electrons is a stronger
attraction compared to one proton to one electron of a
hydrogen atom.
Molecular Compounds
We can draw Lewis diagrams showing the
arrangement of valence electrons in covalent
compounds. In these diagrams we represent each pair
of electrons between atoms as a line.
So for the H2 molecule discussed previously the Lewis
diagram would be:
H–H
All other electrons are represented by dots as
described previously.
Lewis Structures
Draw Lewis Structures of the following molecular
compounds
H
H Nonbonding
electons
a. H2O
O
H
H O
Note each element has a
Noble gas structure by
electron sharing
b. NH3
H
N
H
HN H
H
H
Covalent bonding e’s
Simplified Lewis Structures
Straight lines are used to indicate a shared pair, or
a covalent bond.
H O
H
Nonbonding electrons
Lewis Structure Construction
Step 1
Step 2
Step 3
Step 4
Step 5
Connect each element with a single line
Use the “P” formula to determine extra bonds
Insert the extra bonds, to make double or triple bonds.
Give each atom an octet of electrons, except hydrogen
Determine the formal charge of each element
P = 8(n-q) +2q - 2(n-1) - v
N = number of atoms in molecule
Q = number of hydrogen atoms
V = total number of valence electrons
Examples: Give Lewis Structures for the following
CO2
H2CO3
SO3
NO2+
Lewis Structure of Carbon Dioxide
First, connect atoms with lines
O C O
Lewis Structure of Carbon Dioxide
First, connect atoms with lines
O C O
Second, use “p” formula to determine the number of
extra bonds.
P = 8(n-q) + 2q – 2(n-1) - v
P = 8(3-0) + 2(0) – 2(3-1) - 16
P = 24 + 0 – 4 - 16
P=4
4 extra bonding electrons
2 extra bonds
2 extra lines
Lewis Structure of Carbon Dioxide
Third, add extra lines (Three possible locations)
O
C
O
O
C
O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines (three possible locations)
O C O
O C O O C O
Fourth, give each atom an octet of electrons
Lewis Structure of Carbon Dioxide
Third, add extra lines (three possible locations)
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Lewis Structure of Carbon Dioxide
Third, add extra lines (three possible locations)
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
C O
O C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
-
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
O C O
C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
-
+
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
O C O
C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
-
-
+
O
C
O
O C O
O
C
O
Lewis Structure of Carbon Dioxide
Third, add extra lines
O C O
O C O O C O
Fourth, give each atom an octet of electrons
O
O C O
C O
O C O
Fifth, give each an atom a formal charge
If the element has more than its valence, then it is
negative
If the element owns less than its valence, then it is
positive
-
O
C
O
+
-
+
O C O
O
C
O
Lewis Structure of SO3
There are actually three possible Lewis structures for
SO3.
-
O
S
O
-
O
-
O
O
S
O
-
-
O
S
-
O
O
Each of these three structures is equivalent. We say
they are in “resonance” or that they are “resonance
structures”. Resonance demonstrates how the
loosely held pi electrons are free to move about.
Or another way to look at it is that the pi electrons
are spread over the entire molecule, making it
more stable
Molecular Geometry
So far we have been considering how electrons are
distributed between atoms in molecules and
polyatomic ions.
An important question is:
How can we predict the shape of molecules and
polyatomic ions?
Molecular Shape
The simplest polyatomic ion or molecule is made of
two atoms:
What is the shape of this type of molecule or ion?
The only one way to join two atoms is with a line. All
diatomic molecules and ions have a linear
geometry. Molecular shape is the geometry defined
by the atoms making up the molecule.
Predicting Molecular Shape
Valence shell electron pair repulsion theory
(VSEPR theory) allows us to predict the 3
dimensional shape of molecules and polyatomic ions
with >2 atoms.
VSEPR theory states that electrons in lone pairs
and bonds move as far away from one another as
possible to minimize repulsive interactions.
Repulsion Angles
Predicting Molecular Shape
For a central atom with three electron regions there
are two possibilities.
Predicting Molecular Shape
For a central atom with four electron regions there are
three possibilities.
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
3 regions
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
3 regions
H O H
H N H
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
3 regions
4 regions
H
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is
the number of bonds and electron pairs.
O S O
H O H
H N H
H
3 regions
4 regions
Predicting Molecular Shape
To determine the correct shape of a molecule we must
first begin with the correct Lewis structure.
We then need to determine how many regions of
electron density are around the central atom. This is the
number of bonds and electron pairs.
O S O
H O H
H N H
H
3 regions
4 regions
4 regions
Predicting Molecular Shape
The arrangement of outer atoms is determined by how
many nonbonding pairs there are around the central
atom.
If there are 2 or less electron regions then the
arrangement will always be linear.
The situation becomes more complex if there are more
than 2 electron regions around the central atom.
Predicting Molecular Shape
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
Predicting Molecular Shape
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
Predicting Molecular Shape
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
N
N
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) -10
P = 4 , 2 extra lines, right?
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
N
N
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) -10
P = 4 , 2 extra lines, right?
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
N
N
N
N
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) 10
P = 4 , 2 extra lines, right?
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
N
N
N
N
N
N
Adding nonbonding
pairs
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) -10
P = 4 , 2 extra lines, right?
Adding extra
bonds
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) -10
P = 4 , 2 extra lines, right?
N
N
N
N Adding extra bonds
N
N Adding nonbonding pairs
Consider either nitrogen to
be the central atom. Notice there
are two volumes of space for
electrons
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
First we will start with the Lewis structure of nitrogen.
N
N
N
N
N
N
P = 8(n-q) + 2q -2(n-1) – v
P = 8(2-0) + 2(0) -2(3-1) -10
P = 4 , 2 extra lines, right?
Adding extra bonds
Adding nonbonding pairs
Consider either nitrogen to
be the central atom. Notice there are
two volumes of space for electrons
Two volumes of
electrons repel to 180°,
thus linear shape
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
Now consider the Lewis Structure of water.
H
O
H
Molecular Shape of Water
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and
NH3
Now consider the Lewis Structure of water.
H
O
H
In water there are two bonding
and two nonbonding electron
pairs
Molecular Shape of Water
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and
NH3
H
O
H
Nonbonding
electrons
H
H
In water there are two
bonding and two
nonbonding electron
pairs.
Molecular Shape of Water
Use VSEPR theory to determine as much as possible
about the structure of N2, H2O, SO3, CH4, NH4+ and NH3
H
O
In water there are two bonding
H and nonbonding electron pairs
Nonbonding electrons
Bent or v-shape, bond
angles less than
H
H
109.5°
Electronegativity
Electronegativity is an atoms affinity for
electrons. Affinity is a Greek word for
loving. Since electrons are attracted to the
nucleus, then it makes sense that small
atoms with more protons should attract
electrons stronger. Therefore, we should
predict electronegativity should increase
from left to right and bottom to top on the
Periodic Chart.
Electronegative Chart
Molecule Polarity
• A molecule will be polar if
– it has polar bonds, and
– its centers of partial positive and partial
negative charges lie at different places
within the molecule
• Carbon dioxide, CO2, has two polar bonds
but, because of its geometry, is a nonpolar
molecule
Water
• Water, H2O, has two polar bonds and, because
of its geometry, is a polar molecule
center of partial positive
charge is midway between
the two hydrogen atoms
O
H + H
Water
(a polar molecule)
Ammonia
• Ammonia, NH3, has three polar bonds and,
because of its geometry, is a polar molecule
center of partial
positive charge is
midway between
the three hydrogen
atoms
N
H
H
+
H
Ammonia
(a polar molecule)
Dichloromethane
• Both dichloromethane, CH2Cl2, and
formaldehyde, CH2O, have polar bonds
and are polar molecules
-
-
Cl + Cl
C
HH
D i ch lorome th an e
O-
C+
H
H
Form ald eh yd e
The End
CHAPTER #10