Title: Lesson 2 Covalent Bonding
Learning Objectives:
• Refresh knowledge and understanding of covalent bonding
• Describe the difference between pure covalent and polar covalent
Refresh

Predict and explain which of the following compounds are
ionic:







NaCl
BF3
CaCl2
N2O
P4O6
FeS
CBr4.
Main Menu
Covalent Bonding
= the sharing of PAIRS of electrons between
NON-METAL atoms to form molecules or
giant structures.
i.e. bonding between atoms of SIMILAR and
HIGH electronegativities
A covalent bond is defined as :
“ a shared PAIR of electrons”
 a double covalent bond = 2 shared pairs (4) of e& a triple covalent bond = 3 shared pairs (6) of e-
Covalent Structures

Molecular


As in water and methane
Giant lattice

As in silicon dioxide
Main Menu
Why do covalent bonds form?
Covalent bonds often form
between atoms with too many
electrons in their valence shells to
give away, but not enough to easily
fill.
Thus they share electrons with their
neighbours, in such a way that including the
shared electrons the shells are full
Delocalizing electrons over two atoms instead
of one lowers the energy of the system
COVALENT BONDING
• atoms share electrons to get the nearest noble gas electronic configuration
• some don’t achieve an “octet” as they haven’t got enough electrons
eg
Al in AlCl3
• others share only some - if they share all they will exceed their “octet”
eg
NH3 and H2O
• atoms of elements in the 3rd period onwards can exceed their “octet” if
they wish as they are not restricted to eight electrons in their “outer shell”
eg
PCl5 and SF6
Stopwatch Graph Home
What holds covalently bonded atoms together?
= the electrostatic attraction
between the positive nuclei and
the negative electron pairs shared
between those nuclei.
+
+
Electron Density Maps
Simplified Electron
Density Map of
Hydrogen
This shows the distribution of electrons within the H-H bond / molecule
in terms of contours of equal electron density
The contours which are common to both atoms provide significant
support for the idea of electron sharing in covalent bonding
In contrast, similar maps
for ionic substances show
no overlap
Simplified Electron
Density Map of
Sodium Chloride
Hydrogen H
: 1s1
Chlorine Cl
: 1s22s22p63s23p5
Each atom will share ONE electron from the other
to attain a noble gas configuration
LONE PAIR
a SINGLE
of electrons
covalent bond
Hydrogen
Chloride
H
x
..
.
Cl
..
..
H
..
x
.
Cl
..
A hydrogen chloride
molecule, HCl
..
HYDROGEN CHLORIDE
Cl
H
Chlorine atom
needs one electron
to complete its
outer shell
Hydrogen atom also
needs one electron
to complete its outer
shell
WAYS TO REPRESENT THE MOLECULE
H Cl
H
Cl
PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION
Stopwatch
Graph Home
Oxygen
Oxygen O : 1s22s22p4
To form an oxygen molecule, O2, each atom will share
TWO electrons from the other  noble gas configuration
x
x
O
x
x
x
x
O
x
x
x
a DOUBLE
covalent bond
x
x
O
x
x
x
x
X
X
X
X
x
An oxygen
molecule, O2
x
x
O
x
x
OXYGEN
O
O
O
each atom needs two electrons
to complete its outer shell
O
each oxygen shares 2 of its
electrons to form a
DOUBLE COVALENT BOND
O
O
Stopwatch Graph Home
These dot-and-cross diagrams can be abbreviated :
H
O
Cl
O
Represents a
covalent bond
i.e. a shared
pair of
electrons
SIMPLE MOLECULES
Orbital theory
Covalent bonds are formed when orbitals, each containing one electron,
overlap. This forms a region in space where an electron pair can be found;
new molecular orbitals are formed.
orbital
containing 1
electron
orbital
containing 1
electron
overlap of orbitals provides
a region in space which can
contain a pair of electrons
The greater the overlap the stronger the bond.
Stopwatch Graph Home
Hydrogen H
: 1s1
Carbon C
: 1s22s22p2
One C atom will share ONE electron from each of 4 H
to attain noble gas configurations
Methane
H
.
H
x4!
x
.
.
H
x
.
x
C
.
.
x
H
x
C
.
.
H
A methane
molecule, CH4
METHANE
H
H
H
C
H
atom needs four
electrons to complete
its outer shell
DOT AND
CROSS
DIAGRAM
C
H
H
each atom needs one
electron to complete
its outer shell
H
H
Carbon shares all 4 of its
electrons to form 4 single
covalent bonds
H
H
H C H
H C H
H
H
Stopwatch Graph Home
Hydrogen H
: 1s1
Nitrogen N
: 1s22s22p3
One N atom will share ONE electron from each of 3 H
to attain noble gas configurations
Ammonia
H
.
H
x
.
.
H
x
x3!
LONE PAIR
of electrons
.
x
N
.
..
x
N
.
..
H
An ammonia
molecule, NH3
AMMONIA
H
H
H
N
H
atom needs three
electrons to complete
its outer shell
N
H
H
each atom needs one
electron to complete
its outer shell
Nitrogen can only share 3 of
its 5 electrons otherwise it will
exceed the maximum of 8
A LONE PAIR REMAINS
H N H
H
H
N
H
H
Stopwatch Graph Home
NB : The number of covalent bonds to any
particular non-metal atom is usually
8 (–) it’s group number
GROUP
IV
V
VI
Examples
C , Si
N,P
O,S
NUMBER OF
BONDS PER
ATOM
4
3
2
VII
0
H, F,
He,Ne,
Cl, Br,
Ar
I
1
0
EXAMPLES :
Draw dot-and-cross diagrams to represent the covalent
bonding in each of the following molecules :
1
2
3
4
5
6
7
8
Hydrogen
Nitrogen
Water
Carbon dioxide
Hydrogen peroxide
Tetrachloromethane
Ethane
Ethene
Click Here for Answers
H2
N2
H2O
CO2
H2O2
CCl4
C2H6
C2H4
1
H2
H
H
2
N2
N
N
3
H2O
H
O
H
4
CO2
O
C
O
ANSWERS 5-8
H
O
H
O
H
H
C
C
H
H
7 C2H6
H
5 H2O2
H
8 C2H4
6 CCl4
Cl
Cl
C
Cl
C
C
Cl
BACK TO ANSWERS 1-4
BACK TO SLIDESHOW
Formation of the bond stabilises the atoms
The forces of attraction between
the nuclei and shared electrons are
balanced by the forces of repulsion
between the two nuclei, and this
holds the atoms a fixed distance
apart.
What are the characteristics of covalent substances?
1. ALL covalent substances are non-conducting
because no mobile ions or electrons present (except graphite!!)
2. SIMPLE covalent molecules are :
(a) Gases eg HCl, O2, N2, H2, CH4 etc
(b) Liquids eg H2O, Br2, CH3CH2OH etc
or (c) Low melting point solids eg I2
Strong covalent bonds within the molecules
DO NOT BREAK during melting or boiling
Weaker intermolecular forces between
molecules DO BREAK during melting or boiling
 less energy needed
 lower melting and boiling points
ALL have
LOW melting
& boiling
points
3. GIANT covalent structures are ALL high melting
point solids because a lot of energy is needed to
break the STRONG covalent bonds
Strong covalent bonds can exist throughout
most, or all of, the structure
Most
All
Diamond
Weak Van
der Waal
forces
between
layers
Graphite
SIMPLE COVALENT MOLECULES
Bonding
Atoms are joined together within the molecule by covalent bonds.
Electrical
Don’t conduct electricity as they have no mobile ions or electrons
Solubility
Tend to be more soluble in organic solvents than in water;
some are hydrolysed
Boiling point
Low -
e.g.
intermolecular forces (van der Waals’ forces) are weak;
they increase as molecules get a larger surface area
CH4
-161°C
C2H6
- 88°C
C3H8
-42°C
as the intermolecular forces are weak, little energy is required to
to separate molecules from each other so boiling points are low
some boiling points are higher than expected for a given mass
because you can get additional forces of attraction
Stopwatch Graph Home
Short bonds are strong bonds
Every covalent bond is characterised by two values. (Data will be
given in the IB data booklet.
•Bond length: a measure of the distance between two bonded nuclei.
•Bond strength: usually described in terms of bond enthalpy – a
measure of the energy required to break the bond.
As atomic radius increases down the group the shared electron pair will be further
from the pull of the nuclei, so the bond is weaker.
Multiple bonds have a greater number of shared electrons so have a stronger
force of attraction between the bonded nuclei.
•Greater pulling power = closer together
•Bonds are shorter and stronger
We can also compare two
different carbon-oxygen
bonds within the molecule
ethanoic acid, CH3COOH
Introduction to polar bonds
GAS
MELTING POINT /K
N2
63
CO
74
Both non-metal + non-metal
 expect COVALENT bonding
Both gaseous
 expect SIMPLE MOLECULAR structure
 N2 and CO have SAME Mr (28)
 expect SAME INTERMOLECULAR FORCES
 expect SAME MELTING POINT
WHY NOT?
Explained by BOND POLARISATION
POLARISATION OF COVALENT BONDS
This depends on different atoms having
slightly different ELECTRONEGATIVITY
DEFINITION:
Learn!
The ELECTRONEGATIVITY of an atom is a
measure of its tendency to ATTRACT the
electron pair(s) from a covalent bond
If combining atoms’ electronegativities are :
(a) Very DIFFERENT
eg Na and Cl
 ionic bonding
Na+Cl-
(b) SIMILAR
eg C and H
 covalent bonding
CH4
Pauling Electronegativity Index is used to measure
electronegativity.
H
2.1
Li
1.0
Be
1.5
TRANSITION
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
ELEMENTS
Al
1.5
Si
2.1
P
2.1
S
2.5
Cl
3.0
K
0.8
Ca
1.1
ALL SIMILAR
Ga
1.6
Ge
1.9
As
2.0
Se
2.4
Br
2.8
Rb
0.7
Sr
1.0
IN RANGE
Sn
1.8
Cs
0.6
Ba
0.9
1.5 - 2.0
Pb
1.7
I
2.5
Increases across periods because the number of protons (+)
increases causing increased attraction on the bond electrons.
Decreases down groups because the bond electrons become more
distant and more shielded from the attracting protons
For a “normal” covalent bond
EQUAL sharing of the electron pair occurs
This occurs for atoms with similar electronegativities
Such bonds are called NON-POLAR bonds. (Can
be referred to as ‘pure covalent’ bonds
Examples
H-H
F-F
Cl-Cl
O=O
NN
C-H
Most bonds are POLAR !!!!!
Polar Covalent Bonding
Occurs when electron pairs are shared UNEQUALLY
This occurs for atoms of dissimilar electronegativity
 Bond electrons NOT centralised between the atoms
 e- pair pulled towards more electronegative atom (Y)
and away from less electronegative atom (X).
 slight negative charge (-) on Y,
and slight positive charge (+) on X.
Represented by
+ X  Y -
- called a “dipole”
NB The molecule is still NEUTRAL overall because
+ and - cancel.
Insert the dipole on each of the following:
+ H-Cl -
-
+
-
I - Br -
F-Br +
O-S +
+
+
C-O C-F -
‘Dipole’ is often used to indicate the fact that this type of
bond has two separated opposite electric charges.
Polar bonds  more IONIC character
 stronger attractions between
neighbouring molecules
+ H-Cl -
+ H-Cl -
 called DIPOLE-DIPOLE intermolecular forces
 higher m pt / b pt than expected
N
+ C
N
O -
N
+ C
N
O -
N2 and CO have
same Mr (28) but
mpt CO > mpt N2
because of
polarity of CO.
Electron pair is pulled more strongly
by the Cl atom = polar molecule
Ways to denote the polar nature of a bond
in a structure:
•Write the partial charges over less or
more electronegative atom respectively.
•Write a vector with an arrow indicating
the pull on the electrons by the more
electronegative atoms.
Can you see how they are both used on the
diagram of water?
Pauling Scale of Values
These values will be given in the IB Data Booklet.
Electronegativity is a measure of the ability of an atom to attract electrons in a
covalent bond.
Polar bonds are intermediates
Polar bonds are intermediates between IONIC BONDS and PURE
COVALENT BONDS.