REMEMBER
the
BONDS?
and REMEMBER
the
BON_ _?
….BONES…..
Ionic Bonds: One Big Greedy Thief Dog!
Polar Covalent Bonds: Unevenly
matched, but willing to share.
Metallic Bonds: Mellow dogs with plenty
of bones to go around.
Electronegativity
is
a
quantitative measure of the
ability of an element to
attract the electrons shared
in a bond between it and
another element.
The value of electronegativity goes from
a low of 0.7 for Fr and Cs to a high of 4.0
for F. The higher the value of
electronegativity, the greater the ability
to attract the shared electrons. As a
general trend, electronegativity increases
from left to right across a period and
decreases from top to bottom as you go
down a group.
3 TYPES OF COVALENT BONDING
1. Pure or Non-polar Covalent
Bonding
2. Polar Covalent Bonding
3. Coordinate Covalent Bonding
OUTCOMES
321-4 illustrate and explain the formation of covalent
bonds
Specific Outcomes: At the end of this lesson you should be able
to:
a. Describe the covalent bond as the result of electron sharing.
b. Draw the electron distribution of single and multiple bonds in
molecules .
c. Deduce the Lewis structures of molecules and ions for up to 4
electron pairs on each atom.
COVALENT BOND
a bond formed by the
sharing of electrons
Covalent Bond
A covalent bond is a type of
chemical (JAMES) bond characterized
by the sharing of a pair of electrons
(bones) between two atoms (DOGS).
The electron pair interacts with the
nuclei of both atoms, and this
attractive interaction holds the atoms
together.
Covalent Compounds
Compounds formed between two nonmetals.
 The atoms share the electrons.
• The properties of a molecular
substance—a substance that has atoms
held together by covalent rather than ionic
bonds—are more variable than the
properties of ionic compounds.

Covalent Bond

A covalent bond is a type of chemical
bond characterized by the sharing of a
pair of electrons between two atoms. The
electron pair interacts with the nuclei of
both atoms, and this attractive interaction
holds the atoms together.
Covalent Bond
Between nonmetallic elements of similar
electronegativity.
 Formed by sharing electron pairs
 Stable non-ionizing particles, they are not
conductors at any state
 Examples; O2, CO2, C2H6, H2O, SiC

Covalent Bond
Bonds in all the
polyatomic ions and
diatomic are all
covalent bonds
Pure or Non-Polar
Covalent Bond
* Electrons are equally
shared
Pure or Non-Polar Covalent Bond
• Two of the same atoms form a covalent
bond, for example, when two fluorine
atoms form the fluorine molecule, F2.
• Here is the electron dot structure for F2.
Pure or Non-Polar Covalent Bond
• All other diatomic elements (Cl2, Br2, I2, O2, N2,
and H2) have pure covalent bonds.
• In all these molecules, the electrons are shared
equally.
• Most of the elemental diatomic molecules are
gases at room temperature—Cl2, F2, O2, N2,
and H2.
Pure or Non-Polar Covalent Bond
• Because two atoms of the same
element are forming the bond, the
difference in electronegativities is zero.
• In the fluorine molecule, a pair of
valence electrons are shared equally.
Pure or Non-Polar Covalent Bond
• Seven nonmetal elements are found
naturally as molecular elements of two
identical atoms.
• The elements whose natural state is
diatomic are:
• Hydrogen, Nitrogen, Oxygen, Fluorine,
Chlorine, Bromine, and Iodine
Pure or Non-Polar Covalent Bond
A molecule may be nonpolar either because
there is (almost) no polarity in the bonds
(when there is an equal sharing of electrons
between two different atoms) or because of
the symmetrical arrangement of polar bonds.
Examples of household nonpolar compounds
include fats, oil, and petrol/gasoline. Therefore
(per the "oil and water" rule of thumb), most
nonpolar molecules are water-insoluble
(hydrophobic) at room temperature.
Pure or Non-Polar Covalent Bond
• If two chlorine atoms combine, they
share a single pair of electrons, and each
atom attains a stable octet configuration.
Pure or Non-Polar Covalent Bond
Two oxygen atoms share two pairs of
electrons to form O2, and two nitrogen
atoms share three pairs of electrons to
form N2.
Polar Covalent
Bond
*
Electrons
are
unequally shared
Polar Covalent Bond
• The bond that forms when electrons are
shared unequally is called a polar
covalent bond.
• A polar covalent bond has a significant
degree of ionic character.
Polar Covalent Bond
• Polar covalent bonds are called polar because
the unequal electron
sharing creates two poles
across the bond.
• Just as a car battery or a
flashlight battery has
separate positive and
negative poles, so polar
covalent bonds have poles.
Polar Covalent Bond
• The
negative
pole is centered
on
the
more
electronegative
atom in the bond.
This atom has a
share in an extra
electron.
Polar Covalent Bond
The positive pole is
centered on the less
electronegative
atom. This atom
has lost a share in
one of its electrons.
Polar Covalent Bond
Because there was
not a complete
transfer
of
an
electron,
the
charges on the
poles are not 1+
and 1−, but δ+
and δ−.
Polar Covalent Bond
Because there was
not a complete
transfer of an
electron, the
charges on the poles
are not 1+ and 1−,
but δ+ and δ−.
Polar Covalent Bond
• These symbols,
delta plus and
delta minus,
represent a partial
positive charge
and a partial
negative charge.
Polar Covalent Bond
• This separation of charge, resulting in positively
and negatively charged ends of the
bond, gives the polar
covalent bond a
degree of ionic
character.
Polar Covalent Bond
• The ∆EN for the O — H bond is 1.4, so water
molecules have polar bonds.
• When atoms of hydrogen
and oxygen bond by
sharing electrons, the
shared pair of electrons is
attracted toward the more
electronegative oxygen.
Sample Problem
Writing Lewis Structures
There are few different approaches to write
the Lewis structures, but here one of them
is presented. Follow the following steps.
Step 1. Count the total number of valence
electrons for all the atoms based on the
Lewis dot symbols. For polyatomic atomic
anions add number electrons equal to
number of negative charges. For polyatomic
cations, subtract the number of electrons
equal to the total positive number of positive
charges.
Step 2. Count the total number of electrons
needed to obey the octet rule (8 electrons in
the valence shell) except H that obeys duet
rule ( 2 electrons in the valence shell).
Step 3. Subtract total number of valence
electrons (step 1) from the total number
of electrons needed to obey the
respective rules (step 2) to come with
number of electrons engaged in bonding
( bonding electrons).
Step 4. Divide the bonding electrons
(step 3) with 2 to determine number of
single bond.
Step 5. With information given in either
step 3 or 4 or both, write the Lewis
structure; first write the skeletal
structure using atomic symbols and
making an educated guess which atom
goes in the center and which atoms go
on terminal (end). Most of the times,
you can guess this by looking at the
given molecular formula.
Examples (molecule)
Write the Lewis structure for water (H2O), in which hydrogen atoms are bonded
to the oxygen atom.
Answer:
Lewis dot symbols:
Step 1. Referring to above dot symbols, count the total number of valence
electrons. Total number of valence electrons = 1 (for H) + 1 (for H) + 6 (for O)
= 8 electrons
Step 2: H obeys the duet rule and O obeys the octet rule. Number of electrons
needed to obey the rules = 2 (for H) + 2 (for H) + 8 (O atoms) = 12 electrons.
Step 3. Number of electrons engaged in bonding = 12 electrons - 8 electrons =
4 electrons.
Step 4. Number of single bonds = 4 electrons / 2 electrons = 2 single bonds.
Step 5. Now place the O atom in the center and one H atom on one side and
another H atom on the other side.
Try it!
Try it!