Introduction to Bonding
Topic #13
Essential Question: What are all the differences
and similarities between covalent bonds and
ionic bonds?
What is a bond?
• A bond can be thought of as a force that
holds groups of two or more atoms
together and makes them function as a
unit
• Example : water
O
H
H
Bonds require energy to break
and release energy when made
Types of bonds
• Ionic bonds - typically formed between
metals and nonmetals
• Covalent bonds - typically formed between
nonmetals
• Metallic bonds - formed between metals
Ionic Bonds
• Ionic bonding results from the transfer of
electrons. Then, the opposite charges
attract each other.
• Ionic bonds are strong
Ionic Bonds
• Na and Cl
– Na is a metal and likes to lose one electron
and form a +1 ion.
– Cl is a nonmetal and likes to gain one electron
and form a -1 ion.
– the final ionic compounds is NaCl
Na+ + Cl-
NaCl
The electrostatic interaction
keeps them together!
Ionic Bonds
• They do this
to achieve an
octet!
Covalent Bonds
• Covalent Bonds
– exist between nonmetals bonded together
– form when atoms of nonmetals share
electrons
– electrons can be shared equally or unequally
– The unequal sharing results in polar
molecules
Metallic Bonds
• Metallic bonds exist
between metals
• Occur when two metals,
usually the same metal,
are bonded together
• “sea of electrons”
• “delocalized electrons”
•
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Regents Chemistry
• Electronegativity
How can we tell really tell which
type of bond we have?
• Electronegativity – is the relative ability
of an atom in a molecule to attract
shared electrons to itself
• This tells us what type of bond we have;
– Covalent, polar covalent or ionic
• Electronegativity values are determined
by measuring the polarities of bonds
between various elements to determine
a specific value for each element
Electronegativity
• Electronegativity values for each element
are obtained by using the Periodic Table
• In fact, there is a general trend in
electronegativity we observe in the
Periodic Table
• Electronegativity values increase across
and up the Periodic Table
– See table on pg. 332
Electronegativity
• We take the difference between the
electronegativity values to determine
exactly what type of bond exists, in
essence the polarity of the bond
See table 12.1
Determining Bond Polarity
• If the difference between the
electronegativity values is:
– 0.0 – 0.5: covalent bond (equal sharing)
– 0.6 – 1.6: polar covalent bond (unequal
sharing)
– 1.7 – up: ionic bond (transferring electrons)
Examples
• Use your Reference Tables to determine
the difference in electronegativity values
and the type of bond for each of the
following:
– H-H
– H-Cl
– H-O
– H-S
• H-F
• NaCl
• O2
• KBr
Worksheet
Regents Chemistry
• Intro to valence electrons
Electrons in an atom
• Electrons surround the nucleus of an atom
in specific energy levels or shells
• Each level can hold only a certain amount
of electrons
• It is an atoms ability to the lose, gain or
share electrons from its outer shell that
determine its reactivity
The outer shell
• The outer shell in an atom contains the
valence electrons
• Valence electrons can be lost, gained or
shared to have eight electrons in the outer
shell
• Each group on the table tells the number of
valence electrons
Periodic Table
• Groups 1, 2, 13, 14, 15, 16, 17, 18 have
1,2,3,4,5,6,7,8 valence electrons,
respectively
• We will not consider the transition metals
• See periodic table
Sharing to reach the Octet Rule
• The octet rule states that an atom cannot
have more than 8 electrons in its outer
shell
• Valence electrons are lost, gained or
shared with other atoms to attain 8
electrons in the outer shell
• Eight valence electrons means a filled and
happy shell - like the noble gases
Nonmetals share
• Nonmetals share electrons to reach eight
valence electrons
• Single, double and triple bonds can be
formed by sharing electrons
Metals + non-metals =
lose/gain e• metals and nonmetals interact by losing
and gaining electrons to reach 8 electrons
(filled outer shell)
• The oxidation states on the periodic table
represent this desire to move electrons
• ex: K+ want to lose 1 electron to reach
noble gas configuration of eight electrons
Lewis structures: your
assignment
• The reading and problems focus on
drawing Lewis structures
• Lewis structures are a means to represent
bond formation between atoms
• Covalent bonded compounds have
different Lewis structures than Ionic
bonded compounds
Example of a Lewis Structure
C
H
H
H C H
H
CH4
Covalent bonds
Regents Chemistry
• Lewis Structures
Lewis Structures
 The Lewis Structure is a representation of a
molecule that shows how the valence
electrons are arranged among the atoms in
a molecule
 We used dots around the elemental symbol
to represent the valence electrons
C
Single Lewis Structure - Practice
 Draw four lone electrons first (if necessary)
them pair them up
 Draw Lewis Structures for the following
atoms
Na
Be
Al
Br
Lewis Structures for Ionic
Compounds
 For Lewis Structures of ionic bonds the
atoms are not joined but draw next to each
other
example:
KBr
+
K [
Potassium loses an
electron to achieve
the noble gas configuration
of Argon
Br ]
Bromine gains an
electron to achieve
the noble gas configuration
of Krypton
Lewis Structures – Covalent Bonds
Hydrogen forms stable molecules when it shares
two electrons
 Two electrons fill Hydrogen’s valence shell
 Helium does not form bonds because its valence
shell is already filled; it is a noble gas
 Second row non-metals Carbon through Fluorine
from stable molecules when surrounded by eight
electrons – the Octet Rule

Lewis Structures – Covalent Bonds
 Valence electrons in covalent bonds can
either be bonding pairs, if involved directly
in the bond or lone pairs if not involved in
the bond
Writing Lewis Structures - Rules
Obtain the total sum of the valence electrons from
all of the atoms
 Use one pair of electrons to form a bond between
each pair of bound atoms. For convenience, a line
(instead of a pair of dots) can be used to indicate
each pair of bonding electrons
 Arrange the electrons to satisfy the duet rule for
hydrogen and the octet rule for second row
non metals

Lewis Structures – Covalent Bonds
 Examples
PH
3
H
l
H– P –H
••
Step 1) 8 total valence e- total
Step 2) Draw one pair of
electrons per bond
8-6 = 2 left
Step 3) Arrange the remaining
electrons according to
octet rule
H
H P H
Lewis Structures – Covalent Bond
Practice Examples
HBr
..
H:Br:
··
CF4
Worksheet
Regents Chemistry
– Ionic Lewis Structures
– Multiple bonds in Lewis Structures
– Polyatomic ion Lewis Structures and
Resonance
Lewis Structures for Ionic
Compounds
 For Lewis Structures of ionic bonds the atoms
are not joined but draw next to each other
example:
KBr
+
K [
Br ]
Potassium loses an
electron to achieve
the noble gas configuration
of Argon
Bromine gains an
electron to achieve
the noble gas configuration
of Krypton
Examples of Ionic Lewis
Structures
 Draw Lewis Structures for the following:
NaCl
LiBr
KI
Multiple Bonds and Lewis
Structures…review first
 We have seen how to draw Lewis Structures
for molecules with single bonds
• For example
NH3
1. Sum the total
valence e2. Subtract number
of bonding e3. Place remaining
valence e-
8 total valence e3 bonds x 2e- = 6 bonding
2 e- left over
H N H
H
Multiple Bonds
 Between atoms of the same element
 Example
• Oxygen
O O
O=O
Just
O = O
Also a Lewis Structure
is called a structural model
Example of Multiple Bonds
Nitrogen
N N
N
N
We now meet the octet rule!
Multiple Bonds
 Between atoms of different elements
 CO2
O
C
O
O = C = O
We must use double bonds to meet the octet rule!
Lewis Structures for Polyatomic
Ions and Resonance Structures
 Read pg. 344 (bottom) to 349 and answer
questions a-g in example
12.4 (pg. 347) and a-i in the Self Check
exercise 12.4 (pg. 348)