Introduction to Chemical Bonding
Atoms seldom
 WHY? Because as
exist as
independent particles, atoms
independent
are at relatively high potential
particles in
energy. Nature, however,
nature. Even the
favors arrangements in which
air you breathe is
potential energy is minimized.
By bonding with each other,
made up of
atoms decrease in potential
molecules held
energy thereby creating more
together by
stable arrangements of matter.
chemical bonds.
Introduction to Chemical Bonding

A CHEMICAL
BOND is a mutual
electrical attraction
between the nuclei
(protons) of one
atom and the
valence electrons of
another atom that
binds the atoms
together.
IONIC BOND: the electrical
attraction between large
numbers of cations and
anions
COVALENT BOND: the
sharing of electron pairs
between two atoms.
METALLIC BOND: the
attraction between metal
atoms and the surrounding
sea of mobile electrons of
metals.
IONIC or COVALENT?

Bonds are rarely purely ionic or purely
covalent and usually fall somewhere in
between.
 Recall that electronegativity is the ability of an
atom to attract electrons. This property is
used to determine if a bond is covalent or
ionic.
 You can calculate if a bond is more covalent or
more ionic by finding the difference between
the bonding elements’ electronegativities.
POLAR vs. NON-POLAR
Covalent bonds can be categorized
into two major types: polar covalent and
non-polar covalent.
 Polar covalent bonds involve an
UNEQUAL sharing of electrons between
the two atoms involved in the bond.
 Non-polar bonds involve an EQUAL
sharing of electrons between the two
atoms involved in the bond.

CALCULATING BOND TYPE
 Determining
the type of bond
formed between two types of
atoms is easy. Just subtract
their two electronegativity
values.
Ex:
and
H (electronegativity value: 2.1)
F (electronegativity value: 4.0)
4.0-2.1=1.9
ELECTRONEGATIVITY DIFFERENCES
If the difference is greater than 1.7, the
bond is ionic.
 If the difference is 1.7 or less but greater
than .3, the bond is polar covalent.
 If the difference is .3 or less, the bond is
non-polar covalent.
 0-.3 < .31 – 1.7 < 1.71- 3.3

NONPOLAR
POLAR
IONIC
SECTION REVIEW





What is the main distinction between ionic and
covalent bonding?
How is electronegativity used in determining
the ionic or covalent character of a bond
between two elements?
What type of bonding would be expected
between: H and F __________________
Cu and S __________________
I and Br __________________
6.2



Covalent Bonding &
Molecular Compounds
Molecule- a neutral group of atoms
that are held together by covalent
bonds.
Molecular compound - a chemical
compound whose simplest units are
molecules.
Molecular formula - shows the types
and numbers of atoms combined in a
single molecule of a molecular
compound
6.2



Covalent Bonding &
Molecular Compounds (cont.’d)
Diatomic molecule - a molecule
containing only two atoms.
Bond length – the distance between
two bonded atoms at their minimum
potential energy (the average distance
between two bonded atoms)
Bond energy – the energy required to
break a chemical bond and form
neutral isolated atoms
Bond Lengths vs. Bond Energies
Use Table 6-1 Pg. 168
Graph 10 pairs of coordinates
Using bond length vs. bond
Energy on the grid. What is
Relationship between bond
Length and bond energy?
_________________________
_________________________
_________________________
Binary Molecular Compounds &
Their Nomenclature


A binary MOLECULAR compound is
composed of two different kinds of
nonmetal atoms.
Ex. N2O4
Binary molecular compounds are named
by writing the name of the less
electronegative element first and the
more electronegative element second.
Binary Molecular Compounds
& Their Nomenclature
PREFIXES are used in front of each
element’s name to indicate how many
there are in the formula.
The second element’s name is replaced
with the suffix IDE.
N2O4 : dinitrogen tetraoxide
Se3Cl6: triselenium hexachloride
CO: carbon monoxide
**Notice how the prefix “mono” is used on the 2nd element
but not the first.
Remember: Molecular Compounds are composed of two
non-metals whose electronegativity difference is
between .3 and 1.7. That means that they are bonded
covalently and share electrons.
Also: The prefix system of nomenclature is an older
system of nomenclature. A newer system exists and
will be covered later.
PREFIXES:
Mono
Hexa
Common nonmetals with
“IDE” suffix:
Di
Hepta
Oxide
Nitride
Tri
Octa
Sulfide
Fluoride
Tetra
Nona
Chloride
Iodide
Penta
Deca
Bromide
Phosphide
Selenide
Telluride
Practice Set
Write the name or the formula for
the following molecular compounds.
•
•
•
•
•
•
•
•
•
NO ____________
N2O ____________
NO2 ____________
N2O3 ____________
N2O5 ____________
PCl5 ____________
SF6 ____________
P4O10 ____________
SO2 ____________
• sulfur difluoride
• carbon tetraiodide
• disulfur dichloride
• pentaselenium disulfide
• disilicon octafluoride
• trihydrogen monophosphide
The Octet Rule



Chemical compounds tend to form so that
each atom, by gaining, sharing or losing
electrons, has an octet of electrons in its
highest occupied energy level.
There are few exceptions to the rule: Elements
like H and B only have two and three
electrons respectively. Hydrogen can only
form one bond (2 electrons total) and Boron
can form three bonds (6 electrons total).
Lewis Structures
Lewis Structures: formulas in which electron
dot notation is used to represent molecules.
The electrons between
the H and the Cl atoms
are the shared pair of
electrons. The other
three pairs around the
Cl are called lone pairs
or unshared pairs since
they are not involved in
bonding.
The two dots between hydrogen and
chlorine represent a single bond and can be
represented with a single dash in place of
the dot pair.
When only one pair of electrons are shared,
the atoms have formed a SINGLE BOND.
H Cl
The diagram to the left illustrates a
STRUCTURAL FORMULA which indicates the
kind of atoms, the arrangement and the
bonds but NOT THE UNSHARED PAIRS OF
ELECTRONS.
Examples of Lewis structures
Cl2
O2
N2
H2O
NH3
CH4
Independent Practice
1. SO
3. NBr3
5. H3P
2. CF4
4. CH2Cl2
6. N2
Expanded Valence


Some elements can be surrounded by more
than eight electrons when they combine with
HIGHLY ELECTRONEGATIVE elements
like fluorine, oxygen or chlorine. Ex. SF6
Lewis structure
Couper structure