One Page Lesson: Intermolecular Attractive Forces
The fact that the molecules of most organic compounds exist as liquids and solids at room temperature (rather than flying
about freely as gases) is evidence that the molecules are attracted to one another. The familiar solvent water is a liquid at
room temperature because the H2O molecules are adhering so strongly to one another.
The attractive forces between molecules are collectively referred to as “van der Waals” interactions, after the 19th century
Dutch scientist who first studied them. We’ll simply call them “intermolecular attractive forces,” and focus on three
types: 1) dipole-dipole interactions (dipolar forces), 2) hydrogen bonds, and 3) instantaneous dipole-dipole interactions
(abbreviated IDDI, and also called “dispersion forces” or “induced dipole-dipole interactions”).
A dipole-dipole interaction is the attractive force between the partial positive charge of a polar
group on one molecule and the partial negative charge of a polar group on a neighboring
molecule. Another way to say this is that a region of slight excess positive charge on a molecule
is attracted to a complementary region (of slight excess negative charge) on a neighboring
molecule. Only molecules that contain polar groups of atoms will have this charge imbalance (a
dipole). Polar groups in organic molecules typically contain either an O or a N atom. Organic
molecules that are purely “hydrocarbon” (contain only C and H atoms) are nonpolar, so they are
not capable of adhering by dipole-dipole interactions (electronegativities: H = 2.1, C = 2.5 – so
close that the difference is insignificant, and thus negligible).
You should be able to recognize that the molecule on the top right, with a C=O double bond, has
a dipole (with a slight excess of negative charge on the more electronegative O atom). Two of
these polar molecules would be attracted to one another by a dipole-dipole interaction. The
molecule on the bottom right, made purely of C and H atoms, is nonpolar (has no dipole), so it is
not capable of dipole-dipole interaction.
δ−O
H
C H
H
C
δ+
H C H
H
H
H
H C
C H
H
C
H C H
H
Hydrogen bonds are a very special, rather elite type of dipole-dipole interaction. Although called hydrogen “bonds,” it’s
essential to remember that these are still simply a type of attraction between partially-charged groups of atoms on
neighboring molecules. However, the hydrogen bond, ALWAYS indicated by a dotted line (—O—H . . . N—), is an
especially strong attractive force (about 1/10th the strength of a covalent bond). This is because it consists of two
extremely electronegative atoms (two O, or two N, or one O and one N) with a H atom covalently bonded to one of the
two atoms, but strongly attracted to the other (on a neighboring molecule).
—O—H . . . O—
—O—H . . . N—
—N—H . . . O—
—N—H . . . N—
δ+ H
δ−
H
H
C
C
H
H
O
...
Using the example (—O—H N—), the dotted hydrogen bond is formed because the O atom in
the polar —O—H group is so powerfully electronegative that it draws the shared electrons in the
O—H covalent bond toward itself, leaving the positive charge of H’s nucleus almost fully exposed
(electronegativities: H = 2.1, O = 3.5). ). The δ+ of the H proton is extremely attractive to
electrons on neighboring molecules, particularly to a lone pair of electrons on an electronegative O
or N atom. Thus a hydrogen bond always occurs when a δ+ H attached to an O or N atom on one
molecule is interacting with a δ- O or N atom on a neighboring molecule.
You should now be able to recognize that both molecules on the right contain a polar group of
atoms. But because the molecule on the top has a H atom covalently attached to its O atom, its
polar group is capable of hydrogen bonding. As the molecule on the bottom has no H atom
attached its O atom, it is only capable of the weaker dipole-dipole attractive force.
H C
H
H
δ−O
C
δ+
H C H
H
C H
H
H
All molecules are capable of interacting with one another via instantaneous dipole-dipole interactions (IDDI). To
understand the origin of this attraction, you need to recall the “clouds” of electrons (orbitals) that surround the nucleus of
every atom. You should never think of that electron cloud as being static, or frozen in place; rather, think of it as a
dynamic, swirling mist. Because the electrons are constantly moving, at any given place, the electron cloud may be thin
one instant, and thick the next. Wherever the cloud briefly thickens, the negative charge of the electrons hides the positive
charge of the nucleus. And wherever the cloud is briefly thin, the positive charge on the nucleus shines through. Thus,
whenever two molecules closely approach one another, the temporary partial charges caused by their swirling electron
clouds can interact: the partially-revealed positive charge of one nucleus will be momentarily attracted to the partiallythickened negative charge of a denser electron cloud on the neighboring molecule. Thus, for a brief instant, there’s a very
slight and very fleeting dipole, and the two molecules are WEAKLY attracted to one another. All organic molecules are
capable of this extremely weak instantaneous dipole-dipole interaction; however, this is the only type of intermolecular
attractive force available to molecules that are purely hydrocarbon, and thus nonpolar (have no permanent dipole).
C. Graham Brittain
10/9/2007