We explain what Van der Waals forces are and in which cases they manifest themselves. Also, why do they have such a name and their characteristics.
What are the Van der Waals forces?
It is known as Van der Waals forces or Van der Waals interactions to a certain type of attractive or repulsive intermolecular forces, different from those that generate atomic bonds (ionic, metallic or covalent lattice type) or the electrostatic attraction between ions and others molecules.
Before mentioning the different types of Van der Waals forces, it is important to understand what chemical polarity is. Chemical polarity is a property of molecules that tend to separate electrical charges in their structure.It is a property closely related to intermolecular forces (such as those of Van der Waals), with the solubility and with the points of fusion Y boiling. Depending on the polarity, the molecules can be classified into:
- Polar molecules. They are formed by atoms with very different electronegativity. The atom with the highest electronegativity attracts electrons of the bond and is left with a negative charge density on it. On the other hand, the atom with lower electronegativity will have a positive charge density on it. This distribution of charges will finally lead to the formation of a dipole (system of two charges of opposite sign and equal magnitude).
- Nonpolar molecules. They are made up of atoms with equal electronegativity, so all atoms attract the electrons of the bond in the same way.
A factor that also determines the polarity of a molecule is molecular symmetry. There are molecules made up of atoms of different electronegativity, but which are not polar. This occurs because when the different charge densities of the parts of the molecule are added, they cancel out and result in a null dipole moment.
So, the forces of Van der Waals manifest themselves in three particular ways:
- Keesom attractive forces (dipole-dipole interactions). They are interactions between polar molecules, that is, permanently polarized. Thus, these molecules have a positive pole (with a positive charge density 𝛅 +) and a negative pole (with a negative charge density 𝛅–), and they are oriented so that the positive pole approaches the negative pole.
- Debye attractive forces (permanent dipole-induced dipole interactions). They take place between a polar molecule and an apolar one, but that presents an induced polarity. In this type of interaction, the dipole induces a transient dipole in the apolar molecule.
- London scattering forces (induced dipole-induced dipole). They are interactions that occur between apolar molecules. The movement of electrons in these molecules induces transient dipoles, which causes some attraction between them. They are very weak interactions.
All of these intermolecular forces are known as Van der Waals forces, a name that pays homage to the Dutch physicist Johannes Diderik van der Waals (1837-1923), the first to propose their effects in the equations of state of a gas ( known as the Van der Waals Equation) in 1873. For this discovery he was awarded the Nobel Prize in Physics in 1910.
Characteristics of the Van der Waals Forces
Van der Waals forces grow with the length of the nonpolar end of a substance.
Van der Waals forces are generally weak in comparison with the chemical links ordinary, which does not prevent them from being fundamental to various fields of physical, the biology and engineering. Thanks to them many chemical compounds can be defined.
Van der Waals' forces grow with the length of the nonpolar end of a substance, since they are caused by correlations between fluctuating polarizations between nearby atoms, molecules or surfaces, a consequence of quantum dynamics.
They show anisotropy, that is, their properties vary depending on the orientation of the molecules: it often depends on whether they are attractive or repulsive.
These forces are the weakest that occur between molecules in the nature: Only 0.1 to 35 kJ / mol of energy is required to overcome them. However, they are crucial for the formation of protein.