Van der Waals forces are the sum of the attractive or repulsive forces between molecules. These intermolecular forces, first described by Dutch scientist Johannes Diderik van der Waals, are responsible for many physical properties of matter, such as why liquids turn into solids at low enough temperatures. While individually weak, their collective effect is substantial in chemistry, biology, and materials science.
The Origin of Attraction
At any given moment, the electrons orbiting an atom’s nucleus are in constant, random motion. This movement can lead to a brief, uneven distribution of electrons, creating a temporary, lopsided electrical charge. For a fleeting instant, one side of the atom becomes slightly negative while the opposite side becomes slightly positive. This transient state is known as an instantaneous dipole.
The formation of an instantaneous dipole in one atom or molecule can influence its neighbors. The temporary positive end of one molecule will attract the electrons in an adjacent molecule. This creates a corresponding, induced dipole in the neighboring molecule. This synchronized fluctuation of electron clouds results in a weak, short-lived electrostatic attraction between the two.
This process of inducing dipoles is continuous and happens across vast numbers of molecules. Although each individual attraction is momentary and faint, the sum of these interactions throughout a substance is significant. It is this cumulative effect that gives rise to the cohesive forces that hold non-polar substances together in liquid and solid states.
Types of Van der Waals Forces
Dipole-Dipole Interactions
Certain molecules have a built-in, permanent separation of charge due to their chemical bonds and geometry. These are known as polar molecules, which consistently have a positive and a negative end. Dipole-dipole interactions are the electrostatic attractions between the permanent positive pole of one molecule and the permanent negative pole of another. These forces are stronger than those arising from temporary dipoles because the charge separation is constant. For these forces to be effective, the polar molecules must be close, arranging themselves so that opposite charges align.
Dipole-Induced Dipole Interactions
An attraction can also occur between a polar molecule and a nonpolar one. When a molecule with a permanent dipole approaches a nonpolar molecule, its electric field can distort the electron cloud of the nonpolar molecule. The positive end of the polar molecule will attract the electrons of the nonpolar molecule, inducing a temporary dipole. This induced dipole is then attracted to the permanent dipole of the polar molecule, creating a dipole-induced dipole interaction. This force is weaker than a dipole-dipole interaction because the induced dipole is temporary.
London Dispersion Forces
The most universal type of Van der Waals force is the London dispersion force, named after physicist Fritz London. These forces exist between all atoms and molecules, regardless of whether they are polar or nonpolar. They are the result of the random fluctuations in electron density that create temporary, instantaneous dipoles, as described earlier. The strength of these forces increases with the size of the atom or molecule, as larger electron clouds are more easily distorted or polarized.
Influence in the Physical World
A classic example of Van der Waals forces is the gecko’s ability to scale smooth surfaces, even glass. This feat is not due to suction or a sticky substance, but to the immense number of microscopic hairs, called setae, on its footpads. The cumulative Van der Waals attractions between the molecules of these hairs and the molecules of the surface create a powerful adhesive force.
These forces also dictate the physical states of matter. For instance, noble gases like helium and xenon are individual, uncharged atoms. The only force causing them to condense into liquids at low temperatures is the London dispersion force. Larger atoms like xenon have a larger, more polarizable electron cloud, resulting in stronger dispersion forces and a higher boiling point (-108.1°C) compared to smaller atoms like helium, which has a much lower boiling point (-268.9°C).
In biology, Van der Waals forces are present in the structure of DNA. While hydrogen bonds are the primary forces holding the two strands of the DNA helix together, Van der Waals interactions between the stacked base pairs provide additional stability. These weak attractions help to maintain the double helix’s precise, compact structure, which is necessary for the storage and transmission of genetic information.
Comparison with Stronger Bonds
To understand the role of Van der Waals forces, it is useful to compare their strength to other chemical bonds. They are significantly weaker than the bonds that hold atoms together within a molecule. A Van der Waals interaction has an energy of about 0.4 to 4 kilojoules per mole (kJ/mol). In contrast, covalent bonds, which involve the sharing of electrons between atoms, are much stronger, ranging from 150 to over 400 kJ/mol. Ionic bonds, which involve the electrostatic attraction between charged ions, are also substantially stronger.
This difference in strength is tied to a fundamental distinction in their function. Covalent and ionic bonds are intramolecular forces; they operate within a molecule to hold its atoms together. Van der Waals forces are primarily intermolecular forces, meaning they act between separate, neighboring molecules. This explains why it is possible to boil water by overcoming the intermolecular forces between H₂O molecules, causing them to separate into a gas, without breaking the strong intramolecular covalent bonds that hold the hydrogen and oxygen atoms together within each water molecule.