Polarity and Intermolecular Forces

Review
 We know how to draw Lewis structures for simple
molecules and polyatomic ions.
 We also know how to predict the 3-D geometry of
these molecules and ions, if we apply the VSEPR
Theory.
 Electronegativity (EN) is an atom's tendency to
attract electrons in chemical bonds.
 EN increases to the right and up on the periodic table,
excluding the noble gases.

Bond Polarity
 When two nonmetal atoms bond, they share
electrons.
 They may or may not share electrons evenly.
 Consider the following molecules:
 F2
 HF

Bond Polarity
 H vs. F
 The EN of H is 2.2.
 The EN of F is 4.0.
 In F2, both atoms pull with equal strength on the
bonding e-.
 The e- are shared evenly between them.
 In HF, the F atom pulls harder than the H atom.
 The e- are drawn more toward the F atom.
H
EN = 2.2
F
EN = 4.0
F F F H

Bond Polarity
 EN is same for both
atoms.
 e- density is spread
evenly around
molecule.
 Highest e- density
occurs between atoms.
 Bond is nonpolar.
 EN is much higher for F
than for H.
 e- density is drawn
toward F side.
 F atom acquires partial
negative charge.
 H atom acquires partial
positive charge.
 Bond is polar.
F F F H

Bond Polarity
 In general, a covalent bond is:
 polar if it occurs between two different atoms.
 nonpolar if it occurs between two identical atoms.

Dipole Moments
 Dipole Moment - a measure of the polarity of a
bond.
 Is often represented by a special arrow.
F
H
Arrow points toward
more EN atom.

Polarity of Diatomic Molecules
 Diatomic Molecules - molecules made of only two
atoms.
 If atoms are the same, molecule is nonpolar.
 If atoms are diff., molecule is polar.
 NOTE: Polar does not mean charged.
 Is Cl2 polar or nonpolar?
 Is CO polar or nonpolar?

Molecules With 3 or More Atoms
 A molecule with 3 or more atoms is:
 Polar if its central atom has lone pairs OR
 If the outer atoms are not all the same.
 Nonpolar if its central atom has no lone pairs AND
 All the outer atoms are identical.

CO2 vs. H2O
 Consider the Lewis structure of CO2:
This molecule is nonpolar.

CO2 vs. H2O
 Consider the Lewis structure of H2O:
This molecule is polar.

CH4 vs. CH3Cl
 Neither CH4 nor CH3Cl has any lone pairs on the
central carbon atom.
 Is CH4 polar or nonpolar?
 Is CH3Cl polar or nonpolar?
C C
H
H
H
H
Cl
H
H
H

“Like Dissolves Like”
 Polar molecules mix with each other.
 Nonpolar molecules mix with each other.
 Polar and nonpolar molecules do not easily mix.

Amphipathic Molecules
 Amphipathic - has a hydrophobic region and a
hydrophilic region.
 Hydrophobic - “water-fearing”
 Nonpolar.
 Hydrophilic - “water-loving”
 Polar or charged.
 Dish detergents contain amphipathic molecules.
 Why?

Grease and Water Don't Mix!

Amphipathic Molecules
Hydrophilic head
Hydrophobic tails

Amphipathic Molecules

Intermolecular Forces
 Intermolecular force - a force between two
molecules that does not result from chemical
bonding.
 Dipole-dipole interaction.
 Hydrogen bonding.
 London force.

Dipole-Dipole Interactions
 Dipole - polar molecule.
 Like magnets, except poles are + and  - instead of N
and S.
 Polar molecules generally have higher melting and
boiling points than similar nonpolar molecules.
 EXAMPLE: O2 (nonpolar) boils at -183ºC.
 EXAMPLE: NO (polar) boils at -152ºC.
 NO has a higher boiling point due to its polarity.
 Still far below the boiling point of any ionic cmpd.

Dipole-Dipole Interactions

Hydrogen Bonding
 Hydrogen bond - a stronger form of dipole-dipole
interaction.
 Occurs in molecules that have H atoms bonded to O, N,
or F atoms.
 The small size of the H atom allows these molecules
to get closer together.
 Closer together = stronger forces.
 EXAMPLE: H2O has a boiling point of 100ºC.
 EXAMPLE: H2S has a boiling point of -60ºC.
 The b.p. of H2O is higher b/c of hydrogen bonding.

Hydrogen Bonding
Boiling Points of Several Compounds
-250
-200
-150
-100
-50
0
50
100
150
H2O H2S H2Se H2Te
Compound
Boiling
Point
(ºC)

London Force (Dispersion)
 London force - attraction between temporary
dipoles.
 e- move randomly around molecules.
 Nonpolar molecules become temporarily polar.
 Allows for very weak attractions between nonpolar
molecules.
 Named for Fritz London.

London Forces
+
-

London Forces
 The more e- a molecule has, the greater its London
forces are.
 Large molecules tend to have higher melting/boiling
points than small molecules.
 London forces apply to all molecules.

London Forces
Boiling Points of the Noble Gases
0
50
100
150
200
250
He Ne Ar Kr X
e Rn
Noble Gas
Boiling
Point
(Kelvins)