Organic Chemistry I

Fig. 2.6a
. The dispersion force is weak in nature, and is the weakest intermolecular force. However, since it
applies to all types of molecules (it is the only intermolecular force for nonpolar molecules), dispersion forces are also
the most fundamental intermolecular force.
Figure 2.6a Instantaneous Dipole and Induced Dipole
The magnitude of dispersion forces depends on two factors:
• The relative
polarizability
of electrons. The simple understanding of polarizability is how easily the electrons get
distorted. For larger atoms, there are more electrons in a larger space, therefore the electrons are more loosely
held and more easily polarized, so the dispersion force is stronger. Generally, the larger the molar mass of the
molecule, the stronger the dispersion force.
• The relative
surface area
of the molecule. Molecules with longer, flatter or cylindrical shapes have a greater
surface area compared to the bulky, branched molecules, and therefore have a stronger dispersion force. Taking
the two constitutional isomers of C
4
H
10
(
section 2.1.2
), butane and isobutane as an example, the dispersion force of
butane is stronger than that of isobutane.
•
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Compounds

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Figure 2.6b Shape and surface area effect on dispersion force
Dipole-Dipole Force
For polar molecules, molecules are attracted to each other because of a permanent dipole, and this type of attractive
force is called a dipole-dipole force. As shown below in the electrostatic potential map of acetone, one end of acetone
has a partial negative charge (red) and the other end has a partial positive charge (blue). The dipole-dipole force is an
attraction force between the positive end of one molecule and the negative end of the neighbouring molecule.
Figure 2.6c Electrostatic potential map of acetone
Hydrogen Bonds
First of all, do not let the name mislead you! Although it is called a “bond”, a hydrogen bond is not a covalent bond, it
is a type of intermolecular force. The hydrogen bond is the force between a H atom that is bonded to O, N or F (atoms
with high electronegativity) and the neighbouring electronegative atom,. It can be shown in a general way as:
2.6 Intermolecular Force and Physical Properties of Organic Compounds | 75

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Figure 2.6d Hydrogen bond
The most common example of hydrogen bonding is for water molecules. Water has two O-H bonds, and both are
available as hydrogen bond donors for neighbouring molecules. This explains the extraordinarily high b.p. of water (100
°C), considering the rather small molar mass of 18.0 g/mol. As a comparison, the methane molecule CH
4
with a similar
size has a b.p. of -167.7 °C.
Figure 2.6e Simplified Diagram of Hydrogen
Bonds between Water Molecules
For organic compounds, hydrogen bonds play important roles in determining the properties of compounds with OH or
NH bonds, for example alcohol (R-OH), carboxylic acid (R-COOH), amine (R-NH
2
) and amide RCONH
2
.
The three major types of intermolecular forces are summarized and compared in
Table 2.6
.
76 | 2.6 Intermolecular Force and Physical Properties of Organic Compounds

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Table 2.6 Summary of the Three Major Intermolecular Forces