Chapter 8
Covalent bonding
1
I. Octet Rule

What is the Octet Rule?

The octet rule states that atoms lose
gain or share electrons in to acquire a
full set of 8 valence electrons
Create a drawing of Mg and Cl, Al and Cl
2
II. Covalent Bonds
A Chemical Bond occurs when electron
are shared.
A molecules are formed from the overlap of orbitals
and sharing of electrons
Hydrogen and Hydrogen
3
Covalent bonds
a. Nonmetals hold onto their valence
electrons.
b. They can’t give away electrons to bond.
Still want noble gas configuration.
c. Get it by sharing valence electrons with
each other.
d. By sharing both atoms get to count the
electrons toward noble gas
configuration.
4
Its all in the distance
a.
Too far no bond
b.
Too close electrons repel
c.
Just right and a molecule is born
A molecules are formed from the overlap
of orbitals and sharing of electrons
5
Covalent clip
Covalent_Bonds.asf
6
III. Molecule

A. a covalently bonded compound.
– 1. Tend to occur between non metals
that are close together on the periodic
table.
–
a. diatomic molecules – occur
naturally in nature
a. this is a more stable arrangement.
H H
F F
Br Br
Cl Cl
7
N N
8
I. Single Covalent Bond
A. A sharing of two valence electrons.
1.Only nonmetals and Hydrogen.
2.Different from an ionic bond
because they actually form molecules.
3. Two specific atoms are joined.
9
How does H2 form?

The nuclei repel
+
10
+
When Atoms Combine to make Molecules
Fig 8-1
11
Atoms contain both positive and negative charges. When they come
Together they arrange themselves so that the attractive forges of opposite
Charges is greater than the repulsive forces of like charges
How does H2 form?
The nuclei repel
 But they are attracted to electrons
 They share the electrons

+
12
+
How to show how they formed
It’s like a jigsaw puzzle.
 I have to tell you what the final formula
is.
 You put the pieces together to end up
with the right formula.
 For example- show how water is formed
with covalent bonds.

13
Water
H
O
14
Each hydrogen has 1 valence
electron
Each hydrogen wants 1 more
The oxygen has 6 valence
electrons
The oxygen wants 2 more
They share to make each other
happy
Water
Put the pieces together
 The first hydrogen is happy
 The oxygen still wants one more

HO
15
Water
The second hydrogen attaches
 Every atom has full energy levels

HO
H
16

use acetate sheets to work on page
244 1-5 to create lewis structures
Use a square of acetate to show each
atom (use vis a vis pens only)
 Show the overlap of orbitals
 Draw the outcome in your notebook-use
structural formula (line to represent pair)
 Circle the shared pairs

17
Lewis structure
18

1. PH3

2. H2S

3. HCl

4. CCl4

5. SiH4
Lewis structure
Use molecular model kit to build
 1. PH3

19

2. H2S

3. HCl

4. CCl4
Covalent Bond Formation


Covalent bond forms by overlap of orbitals.
Two types of bonds
Sigma bond: all single bonds are sigma
bonds
Pi bond: in multiple bonds: the first one
is sigma, all other bonds are pi.
There are
20
Single bonds
Multiple bonds (double and triple only)
Two types of Bonds
21

Sigma bonds from
overlap of orbitals
along the axis
connecting the nuclei
between the atoms

Pi bond (p bond):
perpendicular
overlap of p-orbitals
above and below the
axis connecting the
atoms
Sigma bond: s-s Orbital Overlap
22

23
Pg 247- # 12 a-e
III. Multiple Bonds
A. Sometimes atoms share more than
one pair of valence electrons.
B. A double bond is when atoms share
two pair (4) of electrons.
C. A triple bond is when atoms share
three pair (6) of electrons.
24
Carbon dioxide
CO2 - Carbon is central
atom ( I have to tell you)
 Carbon has 4 valence
electrons
 Wants 4 more
 Oxygen has 6 valence
electrons
 Wants 2 more

C
O
25
Carbon dioxide

Attaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
CO
26
Carbon dioxide

Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2
short
OC O
27
Carbon dioxide

The only solution is to share more
O CO
28
Carbon dioxide

The only solution is to share more
O CO
29
Carbon dioxide

The only solution is to share more
O CO
30
Carbon dioxide

The only solution is to share more
O C O
31
Carbon dioxide

The only solution is to share more
O C O
32
Carbon dioxide

The only solution is to share more
O C O
33
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond

O C O
34
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
35
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
36
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
37
How to draw them
Add up all the valence electrons.
 Count up the total number of electrons
to make all atoms complete octet rule.
 Subtract.
 Divide by 2
 Tells you how many bonds - draw them.
 Fill in the rest of the valence electrons
to fill atoms up.

38
Examples
NH3
 N - has 5 valence electrons
wants 8
 H - has 1 valence electrons
wants 2
 NH3 has 5+3(1) = 8
 NH3 wants 8+3(2) = 14
 (14-8)/2= 3 bonds
 4 atoms with 3 bonds

N
H
39
Examples
Draw in the bonds
 All 8 electrons are accounted for
 Everything is full

H
H NH
40
Examples
HCN C is central atom
 N - has 5 valence electrons wants 8
 C - has 4 valence electrons wants 8
 H - has 1 valence electrons wants 2
 HCN has 5+4+1 = 10
 HCN wants 8+8+2 = 18
 (18-10)/2= 4 bonds
 3 atoms with 4 bonds -will require
multiple bonds - not to H

41
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N

HC N
42
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add

HC N
43
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on N to fill octet

HC N
44
Another way of indicating
bonds
Often use a line to indicate a bond
 Called a structural formula
 Each line is 2 valence electrons

H O H =H O H
45
Structural Examples
C has 8 electrons
because each
line is 2 electrons
 Ditto for N

H C N
H
C O
H
46
Ditto for C here
 Ditto for O

Build the molecules using the molecular
model kit.
 HCN
 H 2O
 NH3
 CH4
 C 2H 2
 PH3
 H 2S
 http://www.youtube.com/watch?v=NYFE
5uslaNo

47

48
http://www.youtube.com/watch?v=NYFE
5uslaNo
IV. Bond strength
A. bond strength is the energy needed to
break a covalent bond
1. Bond length and atom size help
determine bond strength
a. The shorter the bond length
The greater the bond strength
i. single bond – longest length
ii. Double bond – medium length
iii triple bond – shortest length
WHICH HAS
THE GREATEST BOND STRENGTH?
49
B. Energy is needed to create and break a
covalent bond.
1.Energy is released when a covalent bond
forms
2. Bond dissociation energy- is needed to
break a bond
a. always a positive number
i. It takes 159KJ/mol to break F2
Would it take more or less to break N2 - why?
50
V. Chemical reaction energy
A. Endothermic reaction – energy is
needed
1. More energy is needed to break the bond
than is needed to create new bond.
AB + energy  A + B
NH4SCN + Ba(OH) + energy ( freezes to wood
because heat is pulled from the reaction
51
B. Exothermic reaction – energy is
released
1. more energy is released when bonds form
than is needed to break bonds
A + B  AB + energy
CaCl + baking soda and water energy is
released as heat
Hw: pg 247 1-12
52
53
I. Naming compounds
A. Two types
1. Ionic - metal and non metal or
polyatomics.
2. Covalent- we will just learn the
rules for 2 non-metals.
54
Covalent compounds
Two words, with prefixes.
 Prefixes tell you how many.
 mono, di, tri, tetra, penta, hexa, Hepta,
octa, nona, deca
 First element whole name with the
appropriate prefix, except mono.
 Second element, -ide ending with
appropriate prefix.
 Practice

55
Writing Formulas
Two sets of rules, ionic and covalent
 To decide which to use, decide what the
first word is.
 If is a metal or polyatomic use ionic.
 If it is a non-metal use covalent.

56
PREFIXES
Mono
- one
 di - two
Tri- three
Tetra- four
 penta- five
 hexa- six
 Hepta- seven
Octa - eight
 nona - nine
 deca - ten
57
Naming Covalent Compounds
CO2
 CO
 CCl4

N2O4
 XeF6
 N4O4
 P2O10

58
Covalent compounds

The name tells you how to write the
formula
Sulfur dioxide
 diflourine monoxide
 nitrogen trichloride
 diphosphorus pentoxide


59
Work on naming ditto
60
I. Acids
Substances that produce H+ ions when
dissolved in water.
 All acids begin with H.
 Two types of acids:
 Oxyacids
 Non Oxyacids- Binary Acids

61
A. Binary Acids

62
1. Binary Acids( Hydrogen and one
other element
– A. hydro + element name + IC + acid
– Examples
– 1. HCl– 2. HF3. HBr-
B. Oxyacids

1. oxyacids- (hydrogen + oxyanions)
– a. name anion + (ic or ous) + acid
• i. use of ic or ous depends on the number of oxygen
atoms in the oxyanion
Examples:
HNO3- nitric acid
HNO2 –Nitrous acid
H3 PO4-phosphoric acid
H2PO3 phosphorous acid
HC2H3O2 –acetic acid
63

H2SO4
sulfuric acid


H2SO3
sulfurous acid
64
Formulas for acids
65

hydrofluoric acid- HF

carbonic acid- H2CO3

hydrosulfuric acid

phosphorous acid
Hydrates
Some salts trap water crystals when
they form crystals.
 These are hydrates.
 Both the name and the formula needs to
indicate how many water molecules are
trapped.
 In the name we add the word hydrate
with a prefix that tells us how many
water molecules.

66
Hydrates
In the formula you put a dot and then
write the number of molecules.
 Calcium chloride dihydrate =
CaCl22O
 Chromium (III) nitrate hexahydrate =
Cr(NO3)3 6H2O

67
68
I. Resonance
A. When more than one dot diagram with
the same connections are possible.
1. SO2
 Which one is it?
 Does it go back and forth.
 It is a mixture of both,
69
Sulfur Dioxide, SO2
Rules 1-3  O—S —O
OR bring in
right pair
bring in
left pair
•
•
••
••
••
O
S
O
••
••
•
•
O
••
••
••
•
•
O
S
+
—
••
•
•
••
•
•
••
••
O— +S
•
•
O
These equivalent structure
are called:
••
RESONANCE
The proper Lewis structure STRUCTURES.
is a HYBRID of the two.
Each atom has OCTET . .
. . . BUT there is a +1 and -1 formal charge
70
20 Oct 97
Bonding and structure (2)
70
Draw the resonance structure
O3
SO3
SO2
71
72
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

CO
73
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

C O
74
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

C O
75
How do we know
Have to draw the diagram and see what
happens.
 Often happens with polyatomic ions and
acids.


76
Work on drawing lewis structures from
packet
BALLOON ACTIVITY
SHOW -LINEAR –TRIGONAL LANAR
AND TETRAHEDRAL – BALLOONS
 DRAW LEWIS DOT AND DEFINE
LONE AND BONDED PAIRS –SHOW
BALLOONS
 HCL
 H2O
 NH4
 CH4

77
VSEPR
Valence Shell Electron Pair Repulsion.
 Predicts three dimensional geometry of
molecules.
 Name tells you the theory.
 Valence shell - outside electrons.
 Electron Pair repulsion - electron pairs
try to get as far away as possible.
 Can determine the angles of bonds.

78
VSEPR
Based on the number of pairs of
valence electrons both bonded and
unbonded.
 Unbonded pair are called lone pair.
 CH4 - draw the structural formula
 Has 4 (C)+ 4(1) (H)= 8
 wants 8 (c)+ 4(2) (H) = 16
 (16-8)/2 = 4 bonds
 Wants-has /2 = how bonds form

79
VSEPR
Single bonds fill
all atoms.
 There are 4 pairs
of electrons
pushing away.
 The furthest they
can get away is
109.5º.

H
H C H
H
80
4 atoms bonded
Basic shape is
tetrahedral.
 A pyramid with a
triangular base.
 Same shape for
everything with 4
pairs.

H
H
109.5º
C
H
81
H
3 bonded - 1 lone pair
Still basic tetrahedral but you can’t see
the electron pair.
 Shape is called
trigonal pyramidal.

H N H H
H
82
N
H
H
<109.5º
2 bonded - 2 lone pair
Still basic tetrahedral but you can’t see
the 2 lone pair.
 Shape is called
bent.

H O
H
83
O
H
H
<109.5º
3 atoms no lone pair

The farthest you can the electron pair
apart is 120º
H
H
84
C O
3 atoms no lone pair
The farthest you can have the electron
pair apart is 120º.
 Shape is flat and called
trigonal planar.

H
H
H
85
C O
H
C
120º
O
2 atoms no lone pair
With three atoms the farthest they can
get apart is 180º.
 Shape called linear.

180º
O C O
86
87
http://connected.mcgrawhill.com/media/repository/protected_con
tent/COMPOUND/50000118/10/23/mole
cular_structure/sco/main.html?stateCod
e=NJ
 PAGE 262- ONLINE TEXT ACTIVITY

88
ABxEy
Notation
How can geometric
structure in 3d be
predicted given the
numbers of valence
electrons?
89
A represents central atom
B represents surrounding atoms
E represents unshared pairs of
valence electrons on Central
Atom
x and y - subscripts that represent
number of surrounding atoms (B)
and
number of unshared pairs of
valence electrons (E)on Central
Atom
90
ABxEy notation
• a good way to distinguish between
electron pair and molecular geometries
is the ABxEy notation
where:
A - atom whose local geometry is of interest
(typically the CENTRAL ATOM)
Bx - x atoms bonded to A
Ey - y lone pair electrons at A
NH3 is AB3E system  pyramidal
91
20 Oct 97
Bonding and structure (2)
91
Example: Water
H2O
Oxygen Central Atom
Hydrogen surrounding atoms
2 surrounding atoms
2 pairs of unshared valence electrons
Notation for ABxEy would be:
92
Answer: AB2 E2
YOUR TURN…
93
CO2
CH4
NH3
PCl5
SF6
BF3
BrF5
IF4
94
VSEPR
Bond
Underlying
Electron
pairs Angles
2
180°
95
Shape
Linear
3
120°
4
109.5°
Tetrahedral
5
90° &
120°
6
90°
Trigonal
Bipyramidal
Octagonal
Trigonal Planar
Actual shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
2
3
3
4
4
4
96
2
3
2
4
3
2
0
0
1
0
1
2
linear
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
5
5
5
5
97
5
4
3
2
0
1
2
3
trigonal bipyrimidal
See-saw
T-shaped
linear
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
6
6
6
6
6
98
6
5
4
3
2
0
1
2
3
1
Octahedral
Square Pyramidal
Square Planar
T-shaped
linear
Hybridization:
The concept of orbital mixing
The simplest compound formed between carbon and
hydrogen is methane, CH . Therefore, it is proposed
that in CH and other carbon compounds, where the
carbon is singly bonded to four other atoms,
covalent bonding involves the hybridization of 2s and
2p orbitals of the carbon atom.
4
4
Hybridization of CH4
_
99
sp3
HYBRIDIZATION
process in which the valence electrons of
an atom are rearranged to form four new
Identical hybrid orbitals
C2H2
100
sp2
sp2
sp3d
sp3d
sp3d
sp3d
sp3d
sp3d2
sp3d2
sp3d2
sp3d2
sp3d2
sp3d2
101
LESSON 11
Types of Covalent Bonding
Types
 Nonpolar
 Polar
102
Non polar and Polar Bonds
When the atoms in a bond are the
same, the electrons are shared equally.
 This is a nonpolar covalent bond.
 When two different atoms are
connected, the electrons may not be
shared equally.
 This is a polar covalent bond.
 How do we measure how strong the
atoms pull on electrons?

103
Electronegativity
A measure of how strongly the atoms
attract electrons in a bond.
 The bigger the electronegativity
difference the more polar the bond.
 0.0 - 0.3 Covalent nonpolar
 0.3 - 1.7 Covalent polar
 1.7>
Ionic

104
Covalent:
Non-polar
0>0.3
Table 8-1
Representative Electronegativity
Differences
Polar:
0.3 < 1.7
Ionic
> 1.7
Pg 335
105
106
Electronegativity
A. Nonpolar Covalent Bonds
1. equal sharing of electrons
2. difference in electronegativity is
(0 - .3 )
3. examples ( H2
O2
N2
Cl2)
107
B. Polar Covalent Bonds
1. unequal sharing of electrons
a. electrons spend more time by
one of the atoms
b. atoms take on a charge
(+dipole or - dipole)
c. difference in electronegativity is
(.3 - 1.7 )
2. examples (HBr H2S CBr4 )
108
C. Ionic Nature of Covalent Bonds
1. polar covalent bonds have an
ionic nature
a. when ionic nature is greater
than 50% the bond is
considered ionic
1) electronegativity difference
is greater than 1.7
109
Work on polar/nonpolar ditto
 Use electronegativity period table

110
Polar Molecules
Molecules with ends
111
Polar Molecules
Molecules with a positive and a
negative end
 Requires two things to be true
 The molecule must contain polar bonds
This can be determined from differences
in electronegativity.
Symmetry can cancel out the effects of
the polar bonds.
Must determine geometry first.

112
How to show a bond is polar
Isn’t a whole charge just a partial charge
 d+ means a partially positive
 d- means a partially negative

d+
H
d-
Cl
The Cl pulls harder on the electrons
 The electrons spend more time near the Cl

113
Covalent:
Non-polar
0>0.3
Table 8-1
Representative Electronegativity
Differences
Polar:
0.3 < 1.7
Ionic
> 1.7
114
Pg 335
Bond polarity concept map
Bond
Covalent
115
Ionic
Molecular bonds concept map

Create a molecular bond concept map
Molecular Polarity

116
Non- polar molecules
polar-molecules
Types of polarity molecules

117
1. Non polar molecules
– Linear- a—a
– Trigonal planar
– Tetrahedral
– Linear a—b—a
2. Polar molecule
Triangular pyramidal
Bent
Linear- a--b
118
D. Polar and Nonpolar Molecules
1. polar molecules
a. asymmetrical molecules with polar
bonds
b. examples (HCl
NH3
H2S)
2. nonpolar molecules
a. molecules with nonpolar bonds
1) examples (H2
Cl2
N 2)
b. symmetrical molecules with polar
bonds
119
E. Properties of Covalent Compounds
1. low b.p. and m.p. compared to ionic
compounds
2. solid (often soft), liquid or gas at room
temperature
3. may form crystalline solids
a. Examples
1) solid CO2 (dry ice)
2) sucrose C12H22O11 (table sugar)
120
How to show a bond is polar
Isn’t a whole charge just a partial charge
 d+ means a partially positive
 d- means a partially negative

d+
H
d-
Cl
The Cl pulls harder on the electrons
 The electrons spend more time near the Cl

121
Is it polar?
HF
 H2O
 NH3
 CCl4
 CO2

122