Polarity in molecules

• Electronegativity describes
how electrons are shared in a
• Consider the compound HCl
• The electron clouds represent where the
two electrons in the HCl bond spend their
time (sizes of atoms are not being shown)
• The shared electrons spend more time
around Cl than H. In other words Cl is
more electronegative than H.
Electronegativity table
• These numbers are derived from several
factors including EA, IE, atomic radius
• You do not need to understand where the
numbers come from
• You need to know that a high number means
the element has a greater pull on electrons
• You will also need to calculate the difference
between values to estimate the % ionic or %
covalent character of a bond
Electronegativity vs Polarity of bond
Calculating EN differences
• The first step in defining the polarity of a
bond is to calculate electronegativity
difference ( EN)
•  EN = EN large - EN small
• E.g. for NaCl,  EN = 2.9 - 1.0 = 1.9
• Next, estimate from fig 7.12 the % ionic
character: about 65% (60 - 70%)
Q-Give the % ionic character for MgO, CH, HCl
MgO = 3.5 - 1.3 = 2.2 … 80% ionic (75-85)
CH = 2.5 - 2.1 = 0.4 … 7% ionic (5 - 10)
HCl = 2.9 - 2.1 = 0.8 … 20 % ionic (15-25)
Note if % ionic is 20%, then % covalent is 80%
Defining polarity
• For our purposes we will define polarity in the
following fashion: 0-10 % is non-polar, 1050% is polar (covalent), 50%+ is ionic
• This is a crude estimate. In reality, the only
non-polar bond between 2 atoms occurs in
diatomic molecules (O2:  EN = 3.5 - 3.5 = 0)
Q - what is the polarity of the bonds in MgO,
CH, HCl?
MgO = 80% ionic = ionic
CH = 7% ionic = non-polar
HCl = 20 % ionic = polar covalent
Polarity in molecules
• A bond that you calculate to be polar may not
be polar if the molecule is symmetrical
• Imagine a tug-of-war between atoms of the
same strength around a central atom
• The pull in one direction is the bond polarity
or “bond dipole”. The overall/molecular
polarity is also known as “dipole moment”
• If the pull is the same from all directions then
the electrons are not attracted to one atom
over another and the molecule is non-polar
Polarity in molecules: Examples
EN of bonds
Bond polarity
Polarity of the
• Why do some solids dissolve in water
but others do not?
• Why are some substances gases at
room temperature, but others are liquid
or solid?
• What gives metals the ability to conduct
electricity, what makes non-metals
• The answers have to do with …
Intermolecular forces
Intermolecular forces
• There are 2 types of attraction in molecules:
intramolecular bonds & intermolecular forces
• We have already looked at intramolecular bonds
(ionic, polar, non-polar)
• Intermolecular forces (IMF) have to do with the
attraction between molecules (vs. the attraction
between atoms in a molecule)
• IMFs come in six flavours: 1) ionic, 2) dipole dipole, 3) H-bonding, 4) London forces, 5) covalent
(network solids), 6) metallic
Ionic, Dipole - Dipole attractions
• We have seen that molecules
+ –
can have a separation of charge
• This happens in both ionic and
polar bonds (the greater the EN,
the greater the dipoles)
H Cl
• Molecules are attracted to each other in a
compound by these +ve and -ve forces
+ –
+ –
H - bonding
• H-bonding is a special type of dipole - dipole
attraction that is very strong
• It occurs when N, O, or F are bonded to H
Q- Calculate the EN for HCl and H2O
A- HCl = 2.9-2.1 = 0.8, H2O = 3.5-2.1 = 1.4
• The high EN of NH, OH, and HF bonds
cause these to be strong forces (about 5x
stronger than normal dipole-dipole forces)
• They are given a special name (H-bonding)
because compounds containing these bonds
are important in biological systems
London forces
• Non-polar molecules do not have dipoles like polar
molecules. How, then, can non-polar compounds
form solids or liquids?
• London forces are named after Fritz London (also
called van der Waal forces)
• London forces are due to small dipoles that exist in
non-polar molecules
• Because electrons are moving around in atoms
there will be instants when the charge around an
atom is not symmetrical
• The resulting tiny dipoles cause attractions
between atoms/molecules
London forces
Instantaneous dipole:
Induced dipole:
Eventually electrons A dipole forms in one atom
are situated so that or molecule, inducing a
tiny dipoles form
dipole in the other
Testing concepts
1. Which attractions are stronger: intermolecular or
2. How many times stronger is a covalent bond
compared to a dipole-dipole attraction?
3. What evidence is there that nonpolar molecules
attract each other?
4. Which chemical has the weakest intermolecular
forces? Which has the strongest? How can you
5. Suggest some ways that the dipoles in London
forces are different from the dipoles in dipole-dipole
6. A) Which would have a lower boiling point: O2 or F2?
Explain. B) Which would have a lower boiling point:
NO or O2? Explain.
7. Which would you expect to have the higher melting
point (or boiling point): C8H18 or C4H10? Explain.
8. What two factors causes hydrogen bonds to be so
much stronger than typical dipole-dipole bonds?
9. So far we have discussed 4 kinds of intermolecular
forces: ionic, dipole-dipole, hydrogen bonding, and
London forces. What kind(s) of intermolecular
forces are present in the following substances:
a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2
(hint: consider EN and molecular shape/polarity)
Challenge: Ethanol (CH3CH2OH) and dimethyl ether
(CH3OCH3) have the same formula (C2H6O).
Ethanol boils at 78 C, whereas dimethyl ether boils
at -24 C. Explain why the boiling point of the ether
is so much lower than the boiling point of ethanol.
Testing concepts
1. Intramolecular are stronger.
2. A covalent bond is 100x stronger.
3. The molecules gather together as liquids or solids at
low temperatures.
4. Based on boiling points, F2 (-188) has the weakest
forces, H2S has the strongest (-61).
5. London forces
– Are present in all compounds
– Can occur between atoms or molecules
– Are due to electron movement not due to EN
– Are transient in nature (dipole-dipole are more
– London forces are weaker
Testing concepts
6. A) F2 would be lower because it is smaller.
Larger atoms/molecules can have their
electron clouds more easily deformed and
thus have stronger London attractions and
higher melting/boiling points.
B) O2 because it has only London forces. NO
has a small EN, giving it small dipoles.
7. C8H18 would have the higher melting/boiling
point. This is a result of the many more sites
available for London forces to form.
8. 1) a large EN, 2) the small sizes of atoms.
Testing concepts
9. a) NH3: Hydrogen bonding (H + N), London.
b) SF6: London only (it is symmetrical).
c) PCl3: EN=2.9-2.1. Dipole-dipole, London.
d) LiCl: EN=2.9-1.0. Ionic, (London).
e) HBr: EN=2.8-2.1. Dipole-dipole, London.
f) CO2: London only (it is symmetrical)
Challenge: In ethanol, H and O are bonded (the
large EN results in H-bonding). In dimethyl
ether the O is bonded to C (a smaller EN
results in a dipole-dipole attraction rather
than hydrogen bonding).
H – bonding and boiling point
Boiling point
Predicted and actual boiling points
Group 4
Group 5
Group 6
Group 7
• Q – Why are some BP suddenly high at period 2?
Testing concepts
Boiling points increase down a group (as
period increases) for two reasons:
1) EN tends to increase
2) Size increases. A larger size means
greater London forces.
Boiling points are very high for H2O, HF, and
NH3 because these are hydrogen bonds
(high EN), creating large intermolecular