Chemical Bond: __________________________
________________________________________
________________________________________
The type of bonding is determined by the
way the valence e-’s are redistributed.
Molecule: ______________
_______________________
_______________________
e.g. H2O
Diatomic Molecule:
molecules containing
__________ atoms.
E.g. O2 , CO , HF, NO
Octet Rule: where chemical compounds tend to
form, such that ______________________________
_________________________________________
This is done by _______________________________
(becoming an ion or entering a covalent bond)
e.g. Fluorine gas exists as F2 .
F –- F
Each F atom has achieved a stable octet by
_______________of electrons.
(do the same with:
O2 , PF3 )
e.g. HCl
Chlorine has a stable
octet, but H does not.
That’s because there are
_______________________
_______________________
_______________________
i.e. __________________
How many electrons would
it need to fill an octet?
Is that possible?
 When
atoms bond
they ______________
_________________
 In dot notation this
is represented as
two dots between
symbols, one from
each atom
(Do F2 , PF3 , HCl )
The unshared pairs of electrons are also known
as ________________.
The _________________________________________.
(do F2 , & PF3 )
These representations are known as:
Lewis Structures: _____________________________
____________________________________________
____________________________________________
The dots representing the lone pairs can
also be dropped. The new
representation is known as ____________
____________________
A single shared pair is known as
a _____________.
Let’s consider O2 :
(Diagram)
The sharing of ____________________ between 2
atoms is known as a ________________.
Let’s consider N2 :
(Diagram)
The sharing of ____________________ between 2
atoms is known as a ________________
Double & Triple Bonds are also known as:
Multiple Bonds.
We still haven’t explained why carbon can form 4
bonds instead of 2…
___________________
Let’s look at Carbon (6):
(ec, orbital diag., Lewis, & Structural)
It makes sense to assume that Carbon forms 2
covalent bonds.
But when Carbon bonds with other atoms, a special
thing happens.
The 2s & 2p merge together to form an
_____________. Now apply Hund’s rule.
(Diagram)
So now, Carbon has 4 single bonds.
Hybridization also applies to Be, B, & Si.
Let’s review with some examples:
(Central atom is the least EN atom or C)
(e.g’s of NH3 , HCN , C2H6 , C2H4 , C2H2)
Show Lewis structure
Ionic: results from 1 atom giving up
its valence e-’s (cation) &
transferring them to another atom
(anion)
e.g. NaCl
Covalent: _____________________
______________________________
 Most bonds are between these
extremes!
Non-Polar Covalent: the bonding valence e-’s
are ________________by the atoms resulting
in _____________________of electrical
charge. e.g. N2
Polar Covalent: the bonding valence e-’s are more
strongly attracted to the more EN atom resulting in
an ___________________________________. It is still
sharing, not a transfer like in ionic. e.g. CO2
So:
Chemical Bonding
Ionic
Covalent
Polar
Non-Polar
The type of bonding can be determined
simply by the _____________________
______________________________(∆EN) of
the 2 atoms. 6.2 p 177.
E.g.’s:
A H-F molecule has an EN difference of:
4.0 (for F) – 2.1 (for H) =
For Na-Cl the EN difference is:
3.0 (for Cl) – 0.9 (for Na) =
For H-H (H2) the EN difference is:
2.1 (for H) – 2.1 (for H) =
The difference tells you what type of bonding that
is occurring:
> 1.7 =
Ionic
< 0.3 =
Non-Polar Covalent
0.31.7 =
Polar Covalent
or:
EN difference =
0
0.3
1.7
3.3
I--------I--------------------I---------------------------I
nonPolar
Ionic
Polar
(see p.162 Fig.6-2 )
Going back to the previous examples:
H-F
Na-Cl
H-H
∆EN = 1.9
∆EN = 2.1
∆EN = 0
Other examples:
Mg-S
CO2



-forces ______________ molecules.
-weaker than ionic & covalent bonds.
In Polar Covalent (∆EN=0.3-1.7)
The EN difference creates a ____________
_________________________________________
_________
e.g. I-Cl => I---Cl
=> I
Cl
(2.5) (3.0)
+ 0.5 Difference
Dipole
Dipole-Dipole interaction
Intermolecular force
______________________
______________________
______________________
Strong intermolecular
force.
The ______________________
force that has H partially
bonded to an electronegative
atom. E.g’s:
Causes higher than normal
boiling points water is a
liquid instead of a gas @ room
temp.
______________________________
____________________________
e.g. H2O , HCl , HF , H2S …
(do diagrams)
I love water!!! (why?)
London Dispersion forces
Average
shape
Temporary
shift
Effects other
molecules
Non-Polar Molecules
(∆EN= 0 - 0.3)
There is no dipole because
the EN diff is too low. But a
__________________________
__________________________
__________________________
______________, which
effects the next molecule,
and so on.
e.g. Draw Cl2 gas with London
dispersion, aka Van der
Waals forces.
Other e.g’s? O2 , H2 , F2 …