Question

The graph below shows the potential energy change as two helium atoms approach and form LDFs. Now consider the interactions between two molecules that are more complex.

Draw a graph of the potential energy change for this interaction (relative to that of He shown in blue).

helium atoms

more complex

molecules

Explain what assumptions you made in drawing your graph. Please talk specifically about how you chose where to position the energy well.

Unlike He atoms, more complex molecules can assume many different orientations with respect to one another. Draw (and label) what you think would be the most stable (lowest potential energy) and least stable orientations for the two molecules.

Explain the reasoning behind your drawings.

Answer 1

When two atoms of helium approach each other LDFs come into play and a attractive interaction develops. In the case of He the drop in potential energy due to the interaction is quite small, that is, the stabilization due to the interaction, and it does not take much energy to knock the two atoms apart. This energy is delivered by collisions with other He atoms. In fact at atmospheric pressures, Helium is never a solid and liquid He boils at ~4 K (−268.93ºC), only a few degrees above absolute zero or 0 K (−273.15 ºC).31 This means that at all temperatures above ~4 K there is enough kinetic energy in the atoms of the system to disrupt the interactions between He atoms. The weakness of these interactions means that at higher temperatures, above 4 K, helium atoms do not “stick together”. Helium is a gas at temperatures above 4 K.

Now let us contrast the behavior of helium with that of hydrogen (H). As two hydrogen atoms approach one another they form a much more stable interaction, about 1000 times stronger than the He–He London dispersion forces. In an H–H interaction the atoms are held together by the attraction of each nucleus for both electrons. The attractive force is much stronger and as the atoms get closer this leads to a larger drop in potential energy and a minimum for the two interacting hydrogen atoms that is much deeper than that for He–He. Because of its radically different stability the H–H system gets a new name; it is known as molecular hydrogen or H2 and the interaction between the H atoms is known as a covalent bond. In order to separate a hydrogen molecule back into two hydrogen atoms, that is, to break the covalent bond, we have to supply energy.32 This energy can take several forms: for example, energy delivered by molecular collisions with surrounding molecules or by the absorption of light both lead to the breaking of the bond.

Distance between atomic centers Potential Energy

Each H can form only a single covalent bond, leading to the formation of H–H molecules, which are often also written as H2 molecules. These H–H molecules are themselves attracted to one another through LDFs. We can compare energy associated with the H–H covalent bond and the H2 - H2 IMF. To break a H–H covalent bond one needs to heat the system to approximately 5000 K. On the other hand to break the intermolecular forces between separate H2 molecules, the system temperature only needs to rise to ~20 K; above this temperature H2 is a gas. At this temperature the IMFs between individual H2 molecules are not strong enough to resist the kinetic energy of colliding molecules.

It turns out that the strengths of LDFs depend on several factors including shape of the molecule, surface area, and number of electrons. For example the greater the surface areas shared between interacting atoms or molecules the greater the LDFs experienced and the stronger the resulting interaction. Another factor is the ability of the electron cloud to become charged, a property known as polarizability. You can think of polarizability as the floppiness of the electron cloud.