Unit 8: Covalent Bonding
But first …
Stretch your mind back to the last chapter…
What is an ionic bond?
An ionic bond is the force of
attraction between two
oppositely charged ions.
+
-
Which of these are not ionic
compounds?
MgCl2
Al2O3
SCl2
K2S
CO2
SCl2 can’t be an ionic compound
because sulfur and chlorine are
both non-metals: they both need
to gain extra electrons to
become stable.
Remember, for an ionic bond to
form, you need an electron giver
and an electron taker.
Draw dot diagrams to show how
one S atom and 2 Cl atoms can
share electrons so they each get
that octet that all atoms want.
Cl S Cl
In a covalent compound, the shared electrons are counted
by both atoms as part of their octets.
That’s what a covalent bond is:
the bond that results when
atoms share electrons.
Sharing = Cooperating:
When you see the “co” in covalent, think of
cooperating and sharing.
How can you recognize a
covalent compound when you
see it?
Hint: Look at these formulas
again.
MgCl2
Al2O3
SCl2
K2S
CO
Properties of Ionic Bonds
1. Most will dissolve in water
2. Will conduct electricity when liquid
(molten) or when dissolved in water
3. Brittle
4. Solids at room temperature
5. High Melting and Boiling Points
6. Made of positive and negative ions
7. Electrons are localized on ions
Properties of Metallic Compounds
1.
2.
3.
4.
WILL NOT dissolve in water
Will conduct electricity when solid
Malleable and Ductile
Most Solids at well above room
temperature
5. High Melting and Boiling Points
6. Made of Positive Ions and delocalized
electrons
Properties of Covalent Compounds
1. Generally WILL NOT dissolve in water
1. Small molecules with OH will dissolve in water –
called Hydrogen Bonding, we will discuss this later.
2. DO NOT conduct electricity
3. Most have low melting and Boiling Points
4. The larger the molecule the greater the chance
it will be a solid at Room Temp
5. Made of Neutral Nonmetals and localized
shared electrons
Drawing Lewis Structures
Procedure
Step 1: Count the valence electrons for all
the atoms in the compound, and add them
up.
Example: To draw the structure of H2O, you’d
add: H: 1 valence electron each x 2 H’s
O: 6 valence electrons
Total: 8 valence electrons
Procedure
Step 2. Write the symbols for the atoms to
show which atoms are attached to which,
and connect them with a single bond.
Our H2O example:
H–O-H
Logical question at this point…
How do you know which atoms are attached to which?
Hints:
a. Hydrogen can form only one bond, so it’s always on an
end, never in the middle.
b. Formulas are often written to show the order the atoms
go in, so HCN is attached this way:
H–C–N
c. When a central atom has group of other atoms bonded to
it, the central atom is usually written first, so CO3 would
be:
O–C–O
O
Procedure
Step 3: Complete the octets of the atoms bonded
to the central atom. (Remember, of course, that
hydrogen can have only two electrons.) Each
bond you already drew represents two electrons
for each atom it’s connecting.
So, for our H2O example, we’ve already completed
this step.
H–O-H
Procedure
Step 4: Place any leftover electrons on the
central atom, even if doing so results in
more than an octet.
If the central atom has 8 electrons, you’re
done. With water, we are done: the
oxygen atom has 4 unshared electrons,
plus two covalent bonds, which count as 2
electrons each.
Procedure
Step 5: If there are not enough leftover
electrons to give the central atom an octet,
try multiple bonds until it does have an
octet.
Use one or more pairs of unshared electrons
on the side atoms to form an extra pair of
shared electrons.
Let’s try another example
Carbon dioxide, CO2.
Step 1: Count the valence electrons you have to work with.
(4 + 6 + 6 = 16)
Step 2: Attach the atoms in a logical way:
O–C–O
You’ve now used 4 of your 14 available valence electrons,
because each single bond represents 2 shared
electrons. So there are 10 electrons left to use.
CO2 example, continued
Step 3: Complete the octets of the outer
electrons. You’ll now have used all 16
valence electrons you have to work with.
However, the carbon atom only has 4 electrons (two from each
single bond).
CO2 example, continued
Step 4: Since the C atom doesn’t have an
octet, move the unshared electrons on the
oxygen atoms until it does:
O
C
O
Now, all three atoms are surrounded by 8 valence electrons, so we’re done!
Draw a dot diagram showing
how ammonia, NH3, is a
covalent compound.
H N H
H
More practice!
1. PCl3
2. H2
3. HCN
4. phosphate ion
5. SF2
6. CO
7. chlorite ion
New definition:
A molecule consists of two or
more atoms bonded covalently.
Identify these chemicals as
molecules, compounds, neither,
or both:
a. CH4
b. MgS
c. I2
d. CO2
e. He
f. Fe2O3
Hydrogen gas and chlorine gas
both exist as diatomic molecules.
“diatomic” means “2 atoms”
Draw dot diagrams of H2 and Cl2.
H H
Cl Cl
These single bonds are also known
as sigma bonds.
(Both words start with “si”)
Oxygen also exists as a
diatomic molecule:
Draw a dot diagram for O2.
O O
The O2 molecule is an example
of a double bond: two pairs of
electrons are shared.
The second bond in a double
bond is called a pi bond.
N2 is another diatomic molecule.
Use dot diagrams to see what type of bond
is in N2.
N N
Another definition:
Bond length: the distance
between the two bonding nuclei
Single bonds are the longest,
double bonds are in the middle,
triple bonds are the shortest.
Single bond
Double bond
Triple bond
The shorter the bond is, the
stronger it is.
Triple bond is the strongest,
so it’s the hardest to break.
The shorter the bond is, the
stronger it is.
Triple bond is the strongest,
so it’s the hardest to break.
In other words, it takes more energy to break a stronger bond.
One last definition:
Bond dissociation energy: the
energy required to break a specific
covalent bond.
Breaking a bond always
requires energy.