All molecules experience intermolecular attractions because molecules attract each other at moderate distances and repel each other at close range. The attractive forces are collectively called “Van der Waals forces” and include all forces that act between electrically neutral molecules in both liquid and solid gases, and in almost all organic liquids and solids. Solids that are held together by Van der Waals forces usually have lower melting points and are softer than those held together by the stronger metallic bonds.
In a nutshell, these forces (especially hydrogen bonds) are essential for building biological systems – our world, in fact. They’re strong enough to bind molecules together but weak enough to be broken, when necessary. intermolecular forces change huge molecules like enzymes, proteins and DNA into the shapes required for biological activity and are responsible for many properties of molecular compounds, including crystal structures (e. g. the shapes of snowflakes), melting points, boiling points, vaporisation, surface tension and densities. Intermolecular forces are weak in comparison to other forces, but, without them, life would be impossible. For example, water would not condense from vapour into solid or liquid forms if its molecules didn’t attract each other.
This scientific law that recognized how molecules interact with each other is named after Dutch physicist, and Nobel Prize winner (1910), Johannes van der Waals, who first put forward the notion of attraction caused by intermolecular forces in 1873 while developing a theory to explain the properties of real gases. It was a revolutionary idea about particle interaction at the time. Van der Waals’ interest in the subject arose in reply to the treatise of another physicist that considered heat as a phenomenon of motion and the existence of “critical temperatures” in gases. Van der Waals demonstrated the necessity of taking into account the volumes of molecules and their intermolecular forces in establishing the relationship between the pressure, volume and temperature of gases and liquids.
Types of Van der Waals forces
There are three types of van der waals forces: dispersion forces, dipole-dipole forces and hydrogen bonding. The various types were first explained by different scientists at different times. Dipole-dipole interactions, for example, were described by Willem Hendrik Keesom in 1912 and dispersion forces by Fritz London in 1930.)
a. Dispersion forces (or London forces) exist between nonpolar molecules and are the weakest links. More electrons in a molecule or atom means potentially larger electron imbalances and so stronger London forces. For example, chlorine gas (CI) consists of two chlorine atoms. In this bond, the electrons are equally shared and are not dominant on one side of the molecule as in HCl (Hydrogen Chlorine). As a result, there may be a tiny point in time where the electrons happen to be dominant on one side However, this temporary charge disappears as quickly as it appears because the electrons are moving so fast. They are always in motion. These temporary dipoles allow the temporarily negative side of one molecule to attract the temporarily positive side of another, which is the intermolecular force. However, dispersion forces between molecules are usually so weak that they can only just float about everywhere and so exist as a gas.
b. Dipole-dipole forces: When the molecule has a distinctly positive end and a negative end (being polar), the permanent force is referred to as a dipole-dipole attraction. Even though the total charge on a molecule is zero, the nature of chemical bonds is such that the positive and negative charges do not completely overlap in most molecules. For example, the oxygen atoms in CO2 (Carbon dioxide) have lots of electrons (more negative), while the carbon atom in the center has far fewer (more positive). It means that the oxygen atom of one CO2 can be attracted to the carbon of another during very close encounters. Such molecules are said to be polar because they possess a permanent dipole moment.
A good example is the dipole moment of the water molecule. Water molecules in liquid water are attracted to each other by Van der Waals forces (or bonds). Even though the water molecule is electrically neutral, the distribution of charge in each molecule is not even and leads to a dipole moment: “a microscopic separation of the positive and negative charge centers. The polar nature of water molecules allows them to bond to each other in groups and is associated with the high surface tension of water” and its viscosity. One could say that Van der Waals forces hold water in its liquid state until thermal agitation becomes violent enough to break those bonds at 100°C. With cooling, residual electrostatic forces between molecules cause most substances to liquify and eventually solidify.
In another example, a molecule like HCl has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent, in-built dipoles will cause the molecules to attract each other rather more than they otherwise would if they had to rely only on dispersion forces.
It’s important to note that dipole-dipole interactions are not an alternative to dispersion forces – they occur in addition to them because all molecules experience dispersion forces. Molecules that have permanent dipoles will have boiling points rather higher than molecules that have only temporary fluctuating dipoles.
c. Hydrogen bonds are abnormally strong dipole-dipole attractions (exactly the same as dipole-dipole interaction), that involve molecules with -OH, -NH, or FH groups. Hydrogen atoms are very small and when a bonded electronegative atom (like oxygen, nitrogen, or fluorine) pulls electrons away from the hydrogen atom, the positive charge that results is tightly concentrated. However, this bond is extremely strong compared to either dipole-dipole forces like HCl because F N and O are very good at attracting electrons and H is equally good at losing them which makes it extremely imbalanced on one side. This results in an extreme dipole situation, thus termed a hydrogen bond. Essentially, the extremely positive side of the molecule will orient itself with the extremely negative side of another molecule.
Sources
http://www.chemguide.co.uk/atoms/bonding/vdw.html
http://209.85.229.132/search?q=cache:fM2ItiE6Oy8J:www-f1.ijs.si/~rudi/sola/Bajd-Van%2520der%2520Waals.pdf
