Question

Distinguish between all the inter-molecular forces and give an example of each. (Please explain as if you'd teach a child....i'm very poor with this).

Answer 1

Inter molecular forces:

Intermolecular forces are forces of attraction or repulsion which act between neighbouring particles (atoms, molecules, or ions). They are weak compared to the intra-molecular forces, the forces which keep a molecule together. Here, atoms within a molecule are attracted to one another by the sharing of electrons. This is called an intra-molecular force. For example the covalent bond, involving the sharing of electron pairs between atoms is much stronger than these forces present between the neighboring molecules.

Attractive intermolecular forces are considered by the following types:

1. Ion-induced dipole forces
2. Ion-dipole forces
3. van der Waals forces (London dispersion force)
4. Hydrogen bonding

Dipole-Dipole Interactions

Many molecules contain bonds that fall between the extremes of ionic and covalent bonds. The difference between the electronegativities of the atoms in these molecules is large enough that the electrons aren't shared equally, and yet small enough that the electrons aren't drawn exclusively to one of the atoms to form positive and negative ions. The bonds in these molecules are said to be polar, because they have positive and negative ends, or poles, and the molecules are often said to have a dipole moment.

HCl molecules, for example, have a dipole moment because the hydrogen atom has a slight positive charge and the chlorine atom has a slight negative charge. Because of the force of attraction between oppositely charged particles, there is a small dipole-dipole force of attraction between adjacent HCl molecules. (refer image 1)

The dipole-dipole interaction in HCl is relatively weak; only 3.3 kJ/mol. (The covalent bonds between the hydrogen and chlorine atoms in HCl are 130 times as strong.) The force of attraction between HCl molecules is so small that hydrogen chloride boils at -85.0oC.

Dipole –Induce dipole forces:

If we mixed HCl with argon, which has no dipole moment, the electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant. (refer Image 2)

Induced Dipole-Induced Dipole Forces:

Neither dipole-dipole nor dipole-induced forces can explain the fact that helium becomes a liquid at temperatures below 4.2 K. By itself, a helium atom is perfectly symmetrical. But movement of the electrons around the nuclei of a pair of neighbouring helium atoms can become synchronized so that each atom simultaneously obtains an induced dipole moment. These fluctuations in electron density occur constantly, creating an induced dipole-induced dipole force of attraction between pairs of atoms. As might be expected, this force is relatively weak in helium -- only 0.076 kJ/mol. But atoms or molecules become more polarizable as they become larger because there are more electrons to be polarized. It has been argued that the primary force of attraction between molecules in solid I2 and in frozen CCl4 is induced dipole-induced dipole attraction.(refer image 3).

van der Waals forces:

The van der Waals forces arise from interaction between uncharged atoms or molecules, leading not only to such phenomena as the cohesion of condensed phases and physical adsorption of gases, but also to a universal force of attraction between macroscopic bodies. Following are the types of van der waals forces,

1. Keesom (permanent-permanent dipoles) interaction
2. Debye (permanent-induced dipoles) force
3. London dispersion force (dipole-induced dipoles interaction)[edit]

London dispersion forces:

The forces that hold molecules together in the liquid, solid and solution phases are quite weak. They are generally called London dispersion forces. Electrons in the orbitals of molecules are free to move around, if you would compare a "snapshots" of a molecule at an instant in time, you would see that there would be slightly different charge distributions caused by the different positions of the electrons in the orbitals. Just how much difference one sees as a function of time is based on the polarizability of the molecule, which is a measure of how well electrons can move about in their orbitals. In general, the polarizability increases as the size of the orbital increases; since the electrons are further out from the nucleus they are less strongly bound and can move about the molecule more easily. Given that two molecules can come close together, these variations in charge can create a situation where one end of a molecule might be slightly negative and the near end of the other molecule could be slightly positive. This would result in a slight attraction of the two molecules (until the charges moved around again) but is responsible for the attractive London dispersion forces all molecules have. However, these London dispersion forces are weak, the weakest of all the intermolecular forces. Their strength increases with increasing total electrons.

Hydrogen bonding:

Hydrogen is a special element. Because it is really just a proton, it turns out that it can form a special type intermolecular interaction called the hydrogen bond. If the hydrogen in a moleucle is bonded to a highly electronegative atom in the second row only (N, O, or F), a hydrogen bond will be formed. In essence the three elements listed above will grab the electrons for itself and leave the hydrogen atom with virtually no electron density (since it had only the one). Now, if another molecule comes along with a lone pair, the hydrogen will try to position itself near that lone pair in order to get some electron density back. This ends up forming a partial bond, which we describe as the hydrogen bond. The strength of this interaction, while not quite as strong as a covalent bond, is the strongest of all the intermolecular forces (except for the ionic bond). (refer image 4)

A hydrogen bond is the attraction between the lone pair of an electronegative atom and a hydrogen atom that is bonded to nitrogen, oxygen, or fluorine. The hydrogen bond is often described as a strong electrostatic dipole-dipole interaction. However, it also has some features of covalent bonding: it is directional, stronger than a van der Waals interaction, produces interatomic distances shorter than the sum of van der Waals radii, and usually involves a limited number of interaction partners, which can be interpreted as a kind of valence.