Prepared by JGL
8/21/2009
IN SUMMARY……
Ionic
bonding
3 types
of
bonding
Metallic
bonding
Covalent
bonding
1
Two hydrogen atoms can share their
valence electrons to attain the same
electron configuration of the nearest Noble
gas configuration, Helium
Prepared by JGL
8/21/2009
COVALENT BONDING
Covalent bond
is the sharing
of two
electrons
between the 2
atoms
2
Covalent Bond
When nonmetallic elements react with other nonmetallic
elements, they share electrons in order to obtain eight valence
electrons.
Each fluorine atom has seven valence electrons. They each
require one more electron to satisfy the Octet Rule.
F F
F F
The left fluorine atom now has a total of eight electrons and the
right fluorine atom now has a total of eight electrons around it.
When nonmetallic elements react with other nonmetallic elements,
they share electrons in order to obtain eight valence electrons.
F F
The two electrons that form the covalent bond are often
Represented by a single line. The F2 molecule can be
represented using a line and dots to show the bonding pair
and the six lone pairs, respectively. This is called a Lewis dot
structure.
F F
Multiple Covalent Bond
Some atoms have to share more than one electron in order
to satisfy the Octet Rule.
O
O
Each oxygen atom has six valence electrons. They each
require two more electrons to satisfy the Octet Rule.
O O
O O
• The left oxygen atom now has a total of eight
electrons around it. The right oxygen atom now
has a total of eight electrons around it.
The four electrons shared by the oxygen atoms form a
double bond.
O O
The double bond is represented by two single lines. Each line
in the Lewis dot structure represents two electrons
The element hydrogen is an exception to the Octet Rule. It
only needs two electrons, rather than eight, to be stable.
H
F
The hydrogen atom has one valence electron. It requires one
more electron to be stable. The fluorine atom has seven
valence electrons. It requires one more to satisfy the Octet
rule.
H F
H F
The hydrogen atom now has a total of two electrons around
it and is stable.
The fluorine atom now has a total of eight electrons around
it and is stable.
H F
H F
• The Lewis dot structure of the HF molecule shows
a line and 6 dots to represent the bonding pair and
the 3 lone pairs of electrons, respectively.
Rules for writing Lewis Dot structures
• Rule 1
Add together the number of valence electrons for each atom
in the molecule. For example, CF4
Carbon has four valence electrons and each fluorine has
seven valence electrons = 4 + 4(7)
= 32
Rule 2
Write out the elements of the molecule so that the least
electronegative elements is in the center surrounded by the
other elements. For example, CF4
F
F C F
F
Rule 3
Place a covalent bond between the central atom and the
outside atoms. Remember each covalent bond contains two
electrons.
F
F C F
F
The four covalent bonds use eight of the 32 valence
electrons in CF4
F
F
C F
• This uses 24 electrons.
There Are no electrons
left, so this is The Lewis
dot structure for CF4
F
• Rule 4
There are 24 valence electrons remaining. Add electrons to
the outer atoms as lose pairs to satisfy the Octet Rule.
Rule 5 for example, NH3
• First apply Rules 1-4 to the molecule
• Rule 1: Count the valence electrons
• Rule 2: Place the least electronegative element at the
centre, except for H which is always an outer atom
• Rule 3: Add covalent bonds between the centre atom and
the outer atoms
• Rule 4: Add lone pairs to the outer atoms
• Rule 5: Add lone pairs to the centre atom
Rule 1
Nitrogen has 5 valence electrons and each hydrogen has 1
valence electron
The total number of valence electrons = 5 + 3 (1) = 8
Rule 2
Hydrogen is always an outer atom and is never at the centre
of a molecule
H NH
H
Rule 3
Add the bonding electrons. This uses 6 of the 8 valence
electrons.
Rule 4
The 2 remaining valence electrons are not added to the outer
atoms, because each H has its maximum of 2 valence
electrons.
• Rule 3
Add the bonding electrons. This uses 6 of the 8
valence electrons.
• Rule 4
The 2 remaining valence electrons are not added to
the outer atoms, because each H has its maximum
of 2 valence electrons.
Rule 5
H NH
H
This is the
Lewis structure
For NH3
Place the remaining 2 Valence
electrons on the central
nitrogen atom
Rule 6
Check all atoms in the molecule to ensure that each has 8
electrons(2 for hydrogen). If an atom has fewer than 8
electrons, create double or triple bonds. (Note: Double
bonds only exist between C,N,O and S atoms)
Apply rule 6 to the following; CH4, CF4,
H
H C H
• Hydrogen : 1 bond = 2
electrons (stable)
• Carbon
: 4 bonds = 8
electrons (stable)
H
F
F
C F
F
• Fluorine : 1 bond + 3
lone pairs = 2 + 3 (2)
= 8 electrons (stable)
• Carbon : 4 bonds = 8
electrons (stable)
Example; CH2O
Apply Rules 1-5 to the molecule
Rule 1: Count the valency electrons
Rule 2: Place the least electronegative element at the
centre, except for H, which is always an outer atom
Rule 3: Add covalent bonds between the centre and the
outer atoms
Rule 4: Add lone pairs to the outer atoms
Rule 5: Add lone pairs to the centre atom
Rule 1
Carbon has 4 valence electrons, each hydrogen has 1 valence
electron, and oxygen has 6 valence electrons.
Total number of valence electrons : 4 + 2(1) + 6 = 12
Rule 2
Carbon is at the centre of the molecule because it is less
electronegative than oxygen. Hydrogen is always an outer
atom and is never at the centre of the molecule.
H
C
O
H
• Rule 3
Add the bonding electrons.
This uses 6 of the 12 valence
electrons
H
C
O
H
H
C
O
H
• Rule 4
Add the remaining 6 lectrons to
the outer atom. Hydrogen does
not need any more electrons, but
Oxygen needs 6 to complete its
octet.
Rule 5 There are no valence electrons left to add to the centre
H
C
H
O
• Rule 6
Oxygen shares one of its
lone pairs with C and O
and give the desired 8
electron total
H
C
H
O
H
This is the
Lewis dot
Structure for
CH2O
C
O
H
Exceptions to the Octet Rule
The Octet Rule applies to Groups IVA through VIIA in the
second row of the Periodic Table, but there are exceptions
to the rule among some other elements. The following two
cases are an example
Example BF3
Rule 1
Boron has 3 valence electrons and each Fluorine has 7
valence electrons
Total number of electrons = 3 + 3 (7) = 24
F
B
F
F
Rule 3
Add the bonding electrons.
This uses 6 of the 24 valence
electrons
Rule 2
Boron is at the centre of
the molecule because it is
less electronegative than
fluorine
F
B
F
F
Rule 4
Add the remaining electrons
to the outer atoms. Each
Fluorine has the required 8
electrons
F
B
F
F
Rule 5
This uses the remaining
electrons leaving none to add
to the Boron central atom
F
B
F
F
Rule 6
Check the number of electrons around each atom. Each
Fluorine atom has 8 electrons, but the Boron Atom has only
6. This is an exception to the Octet Rule. A B=F bond is not
an option, because double bonds exist only between C,N,O,
and S atoms
F
B
F
F
This is the Lewis
dot structure BF3
The electrons are free to move
from one positively charged ION
to the next (i.e. They are
DELOCALISED) and are shared
However,
metals behave
(just like in covalent bonding
differently.
among the various metallic
positively charged ions
Prepared by JGL
8/21/2009
METALLIC BONDING
The valence (outermost)
electrons are loosely held by
the metal ions, so much so
that they move away from
the atom to form a
positively charged ION.
Metallic bonding is similar
to both covalent and ionic
bonding
The number of electrons =
the number of protons.
Source: www.daviddarling.info/images/metallic_bond.jpg
The metal is therefore
electrically NEUTRAL
31
Similarities


Metallic and ionic
bonding involve
electrostatic attractions
between positive and
negatively charged
particles.
Metallic bonding shares
electrons among the
ions in a similar manner
to how electrons are
shared in covalent
bonding.
Differences

Covalent bonding shares
electrons rather than
having electrostatic
charges.
Prepared by JGL
8/21/2009
COMPARE AND CONTRAST TYPES OF BONDING
32