Covalent, Metallic or Ionic Bonds | Chemical Bonds in Molecular Atoms | Salt Crystals | How Atoms and Molecules Move in Liquid? | Lewis Dot Structure and Molecular Shape | Electronegativity, Ionization, and Bond Formation | Van Der Waals Forces
|Van Der Waals Forces
|CCSTD HS Chemistry 2.h.|
Intramolecular bonds exist within molecules and include Covalent, Ionic, and Metallic bonds. These intramolecular bonds are relatively strong.
In this section we will look at intermolecular bonds, which are forces that exist between molecules. Intermolecular bonds, often times called van der Waals Forces, are relatively weak.
Johannes Diderik van der Waals (1837–1923) was a Dutch physicist that who won a Nobel Prize in Physics in 1910 for his work on ideal gases. To explain the non-ideal behavior of gases, van der Waals postulated the existence of weak intermolecular forces, or van der Waals forces.
Although the remainder of this discussion will involve intermolecular forces, we need to first discuss an intramolecular property called the dipole moment.
A dipole moment is the result of two opposing partial electric charges on the same molecule. Technically, the dipole moment is the size of a pair of opposite electrical charges of a molecule times the distance between the charges. A dipole moment (p) is a vector quantity with a magnitude equal to the product of the charge (q) of one of the poles and the distance (d) separating the two poles.
When discussing dipole moments, we assign a positive partial charge (δ+) to one end of the molecule, and a partial negative charge (δ-) to the opposite end.
When a molecule has a dipole moment, then it is called a polar molecule (being as how it has two poles). And of course, a nonpolar molecule has no partial charges, or its partial charges cancel each other out.
A polar bond is partway between a covalent and an ionic bond in
character, although it is much weaker than either. A covalent bond
within a molecule that has a dipole moment is said to be a
Within the classification of van der Waals forces, there are the attractive forces of;
and the repulsive force of the;
First, we will discuss the van der Waals attractive forces.
Some molecules, although electrically neutral, are dipoles and have a relatively positive (δ+) center and a relatively negative (δ-) center. A dipolar molecule has electron densities shifted from the less to the more electronegative atom.
When a molecule with a permanent dipole moment nears another molecule with a permanent dipole moment, those two molecules will have a weak, but significant, attraction. This is the van der Waals attractive force that is called a dipole-dipole interaction (also called a Keesom interaction).
Dipole-dipole interactions between like molecules cause the compound to have greater melting points, boiling points, and viscosities than compounds that have molecules without permanent dipole moments.
Figure 3.7.2 shows the dipole-dipole interaction found between in two
molecules of Hydrochloric Acid. In comparing the strength of the
dipole-dipole interaction with a covalent bond, note that a molecule of
Hydrochloric Acid (HCl) has a Hydrogen-Chlorine bond strength of
430kJ/mol and the intermolecular dipole-dipole interaction is about
Hydrogen bonding is a special case of a dipole-dipole interaction. Let’s examine hydrogen bonding in the water molecule; hydrogen bonding takes place when the electronegative Oxygen atom of a water molecule interacts with one or more Hydrogen atoms from other second water molecules.
The electron density of a water molecule is most heavily shifted toward the Oxygen atom and away from the two Hydrogen atoms. Furthermore, the geometry of the water molecule is ‘bent’ from a linear shape because of the two lone electron pairs on the Oxygen atom.
Figure 3.7.3 shows the arrangement of valence electrons in a water molecule that gives rise to its permanent dipole moment; electrons are shifted to the electronegative Oxygen atom (δ-) and away from the Hydrogen atoms (δ+). The partially positive Hydrogen atoms are free to associate with the partially negative Oxygen atoms of another water molecule.
Figure 3.7.4 demonstrates how hydrogen bonding results in each Oxygen atom of a water molecule having an average of 3.5 ‘bonds’ with nearby Hydrogen atoms. Hydrogen bonding in water is a relatively strong association (~23kJ/mol) and is responsible for many of water’s unique characteristics; i.e. high boiling point, high surface tension, expansion of the solid state, and ability to dissolve ionic and polar molecules.
Hydrogen bonding takes place with molecules other than water; all you
need is a highly electronegative atom on a molecule that includes one or
more a loosely held Hydrogen atoms. For example, hydrogen bonding
contributes to the
tertiary structures of
Want to see an animated graphic of hydrogen bonding in water? Click:
Figure 3.7.5 illustrates the dipole-induced dipole interaction
between a molecule of Hydrochloric Acid, which is a polar molecule with
a permanent dipole moment, and an atom of Argon.
3. London Forces
The weakest of the van der Waals forces are London Forces, also known as dispersion forces, or as instantaneous induced dipole-induced dipole forces.
When electron densities around the two Helium atoms shown in Figure 3.7.6 become synchronized so that dipole moments of the two atoms are instantaneously and simultaneously induced, the atoms will be attracted to each other. The London force between the two Helium atoms is brief and quite weak, having energy of about 0.07 kJ/mol.
Although London forces are the weakest attractive forces of all the van der Waals interactions, they occur constantly, making for fluctuating and transient dipole moments in neutral atoms or in nonpolar molecules.
Because the fluctuations in London forces arise with fluctuating electron densities, those elements further down a column of the Periodic Table have more electrons, and so their fleeting induced partial negative charge (London force) is greater. This is the reason why the noble gases have increased boiling points with increasing size; increasing atomic size means increased London forces.
All of the attractive van der Waals forces become weaker with heating.
In terms of relative strength of the attractive van der Waals forces as compared to an Ionic (or Covalent) bond, we can approximate;
All of the attractive van der Waals forces are diminished with heat.
There is one van der Waals repulsive force;
1. Hydrophobic Effect
The Hydrophobic Effect is the only repulsive van der Waals force. The hydrophobic effect is an exclusion of water to the point that it forces water to bond with its self rather than to bond with a nonpolar molecule, or to the nonpolar portion of a molecule.
The hydrophobic effect is not really an interaction because it actually causes a dipole-dipole interaction (Hydrogen bonding by water) not to take place. The hydrophobic effect is not particularly strong, but is just as important as hydrogen bonding in achieving the tertiary structures (folding behavior) of proteins and of DNA
It is the hydrophobic effect that causes
The hydrophobic effect can be cancelled out by lowering the temperature.
At temperatures near zero (oC), water prefers to be in an ordered
structure and any order generated by the hydrophobic effect is no longer
as energetically favorable.