Covalent Bonds
-Electronegativity
-Polar/non-polar
covalent compounds
-Lewis dot structures
-Resonance
Covalent Bonding – What?
A covalent bond is a form of
chemical bonding that is characterized
by the sharing of pairs of electrons
between atoms, or between atoms
and other covalent bonds.
 Different from Ionic bonding because
the atom does not give up electrons,
or gain them.

What does covalent
mean?

The term covalent bond dates from
1939.The prefix co- means jointly,
associated in action, etc.; thus a “covalent bond”, essentially, means that
the atoms share “valence.” In the
molecule CH4, the hydrogen atoms
share their electrons with those of
carbon via covalent bonding.
Electronegativity
Electronegativity
Covalency is greatest between atoms
of similar electronegativities.
 Covalent bonding does not
necessarily require the two atoms be
of the same elements, only that they
be of comparable electronegativity.

Electronegativity cont.
Electronegativity, symbol χ, is a
chemical property that describes the
ability of an atom to attract electrons
towards itself.
 Electronegativity cannot be directly
measured and must be calculated
from other atomic or molecular
properties.

Pauling Method

There are several ways to calculate
electronegativity but the most
common is to use the Pauling Method
named after the man who first
proposed electronegativity – Linus
Pauling – in 1932
The Pauling Scale

Using the Pauling Method gives
dimensionless quantities placed on
the “Pauling scale”, which runs from
0.7 to 4.0 (hydrogen = 2.2).
Pauling’s method
Electronegativity is calculated by
looking at the difference between two
atoms joined in a covalent bond.
(XA – XB)
 Dissociation energies, Ed are
expressed in Electron Volts (eV)

Example.

The difference in Pauling
electronegativity between hydrogen
and bromine is 0.73
(Ed: H–Br, 3.79 eV; H–H, 4.52 eV; Br–
Br 2.00 eV)
The scale needs a base!
As only differences in electronegativity
are defined, it is necessary to choose
an arbitrary reference point in order to
construct a scale. Hydrogen was
chosen as the reference, as it forms
covalent bonds with a large variety of
elements.
 Electronegativity of hydrogen is 2.20
 Every other electronegativity is based
of hydrogen's. (0.7 – 4.0)

More or less than
Hydrogen?
It is also necessary to decide which of the
two elements is the more electronegative
(equivalent to choosing one of the two
possible signs for the square root).
 This is done by "chemical intuition":
 Example: Hydrogen bromide dissolves in
water to form H+ and Br− ions, so it may
be assumed that bromine is more
electronegative than hydrogen.

How it affects a bond.

Same electronegativity.
If the atoms are equally electronegative, both have the same
tendency to attract the bonding pair of electrons, and so it will
be found on average half way between the two atoms. To get
a bond like this, A and B would usually have to be the same
atom. You will find this sort of bond in, for example, H2 or Cl2
molecules.
Note: It's important to realise that this is an average picture.
The electrons are actually in a molecular orbital, and are
moving around all the time within that orbital.
What if one is more enegative.

B is slightly more electronegative.
That means that the B end of the bond while sharing with A,
has slightly more time with the electrons, so it becomes
slightly negative. At the same time, the A end becomes
slightly positive. In the diagram, " " (read as "delta") means
"slightly“ - so + means "slightly positive". This is also called
a Polar Bond
And if one is a lot more?

B is a lot more electronegative than A
In this case, the electron pair is dragged over to
B's end of the bond. To all intents and purposes,
A has lost control of its electron, and B has
complete control over both electrons. This is now
no longer covalent but an ionic bond.
So when do the bonds
change?


The implication of all this is that there is no clear-cut
division between covalent and ionic bonds. In a pure
covalent bond, the electrons are held on average
exactly half way between the atoms. In a polar bond,
the electrons have been dragged slightly towards one
end.
How far does this dragging have to go before the bond
counts as ionic? There is no real answer to that. You
normally think of sodium chloride as being a typically
ionic solid, but even here the sodium hasn't completely
lost control of its electron. Because of the properties of
sodium chloride, however, we tend to count it as if it
were purely ionic.
Electronegativity
difference predictions
Non-polar covalent bond: < 0.5
 Slightly polar bond: 0.4-0.9
 Moderately polar bond: 1-1.3
 Highly polar bond: 1.4-1.7
 Slightly ionic bond: 1.8-2.2
 Ionic Bond: 2.3+

Summary
No electronegativity difference
between two atoms leads to a pure
non-polar covalent bond.
 A small electronegativity difference
leads to a polar covalent bond.
 A large electronegativity difference
leads to an ionic bond.

Polar and Non-Polar
Compounds
Polar and Non-polar
Polar bonds occur when one atom
attracts electrons more strongly in a
covalent bond.
 This causes atoms to have a slightly
negative and slightly positive “end”

Polar substances are soluble in
water
 An example of a polar
compound is also water - the
electrons of water's hydrogen
atoms are strongly attracted to
the oxygen atom, and are
actually closer to oxygen's
nucleus than to the hydrogen
nuclei; thus, water has a
relatively strong negative
charge in the middle (red
shade), and a positive charge at
the ends (blue shade).

Non-Polar

Non-Polar compounds are formed
when the electrons are shared equally
between two atoms.
Non-Polar
Compounds that are non-polar are
insoluble in water.
 Due to this, there are no positive or
negative “ends” to the compound.
 Examples of Non-polar compounds
include: fats, oil and petrol.

Predicting Molecular Polarity
POLAR
FORMULA
DESCRIPTION
EXAMPLE
XY
Linear molecules
CO
HXz
Molecules with a single H
HCl
XzOH
Molecules with an OH at
one end
C2H5OH
OzX
Molecules with an O at one
end
H2O
NzXz
Molecules with at one end N NH3
Xz
NON
POLAR CzXz
Single molecule
O2
Carbon compounds
CO2
Lewis Structures
G.N. Lewis
Lewis Dot Structures
Lewis Structures were originally created
by Gilbert Newton Lewis, who
introduced them in his 1916 article The
Atom and the Molecule.
 Also called Lewis dot diagrams, and are
similar to electron dot diagrams.
 He developed them to assist him in
teaching chemical bonding, by putting
dots in the place of valence electrons.
Because of this, Lewis structures only
deal with the valence electrons.

Cont.
Atoms always want to have 8
electrons in the outer shell, or in the
case of hydrogen and helium, 2.
 The goal in all bonding is to gain a full
valence shell, or to obtain a noble
structure. This is called the Octet
rule.

Covalent Bonds – Lewis
Diagrams
Here we are only dealing with covalent bonding, all
of the Lewis structures we deal with will be sharing
electrons.
This means that the valence electrons will be
shared between both participating atoms.
This is represented in Lewis structures by either a
line, or a circle encompassing both of the shared
electrons.
Steps to creating a Lewis
Structure
1.
2.
3.
4.
Draw the atoms, and then draw the appropriate
number of valence electrons around the atoms.
Because we are only dealing with covalent
bonding, automatically assume that it is covalent
and pair up the electrons.
Apply the Octet rule.
Determine if it is polar or non polar, and state
which atom is slightly negative, and which atom
is slightly positive. This step is not necessary for
Lewis Structures, but you should be able to do
it.
Resonance
Resonance



Many bonding situations can be described with more
than one valid Lewis Dot Structure
(for example, ozone, O3). In an LDS diagram of O3,
the center atom will have a single bond with one atom
and a double bond with the other. The LDS diagram
cannot tell us which atom has the double bond
These two possible structures are called resonance
structures. In reality, the structure of ozone is a
resonance hybrid between its two possible
resonance structures. Instead of having one double
bond and one single bond, there are actually two 1.5
bonds with approximately three electrons in each at all
times.
Resonance

Lewis Dot Structures for molecules
with resonance are shown by creating
the dot structure for every possible
form, placing brackets around each
structure, and connecting the boxes
with double-headed arrows.