The hydrogen molecule is the most basic example in quantum mechanics of how atoms can combine into molecules in order to share electrons. So, the question suggests itself whether, if hydrogen molecules are brought close together in a solid, will the atoms start sharing their electrons not just with one other atom, but with all surrounding atoms? The answer under normal conditions is no. Metals do that, but hydrogen under normal conditions does not. Hydrogen atoms are very happy when combined in pairs, and have no desire to reach out to further atoms and weaken the strong bond they have already created. Normally hydrogen is a gas, not a metal.
However, if you cool hydrogen way down to 20 K, it will eventually condense into a liquid, and if you cool it down even further to 14 K, it will then freeze into a solid. That solid still consists of hydrogen molecules, so it is called a molecular solid. (Note that solidified noble gases, say frozen neon, are called molecular solids too, even though they are made up of atoms rather than molecules.)
The forces that glue the hydrogen molecules together in the liquid and solid phases are called Van der Waals forces, and more specifically, they are called London forces. (Van der Waals forces are often understood to be all intermolecular forces, not just London forces.) London forces are also the only forces that can glue noble gas atoms together. These forces are weak.
It is exactly because these forces are so weak that hydrogen must be
cooled down so much to condense it into liquid and finally freeze it.
At the time of this writing, that is a significant issue in the
hydrogen economy.
Unless you go to very unusual
temperatures and pressures, hydrogen is a very thin gas, hence
extremely bulky.
Helium is even worse; it must be cooled down to 4 K to condense it into a liquid, and under normal pressure it will not freeze into a solid at all. These two, helium and hydrogen are the worst elements of them all, and the reason is that their atoms are so small. Van der Waals forces increase with size.
To explain why the London forces occur is easy; there are in fact two explanations that can be given. There is a simple, logical, and convincing explanation that can easily be found on the web, and that is also completely wrong. And there is a weird quantum explanation that is also correct, {A.33}.
If you are the audience that this book is primarily intended for, you
may already know the London forces under the guise of the
Lennard-Jones potential. London forces produce an attractive potential
between atoms that is proportional to
(10.1) |
The first term in the Lennard-Jones potential is there to model the
fact that when the atoms get close enough, they rapidly start
repelling instead of attracting each other. (See section 5.10
for more details.) The power 12 is computationally convenient, since
it makes the first term just the square of the second one. However,
theoretically it is not very justifiable. A theoretically more
reasonable repulsion would be one of the form
It may be noted that at very large distances, the London force takes
the Casimir-Polder form
Molecular solids may be held together by other Van der Waals forces besides London forces. Many molecules have an charge distribution that is inherently asymmetrical. If one side is more negative and the other more positive, the molecule is said to have a “dipole strength.” The molecules can arrange themselves so that the negative sides of the molecules are close to the positive sides of neighboring molecules and vice versa, producing attraction. (Even if there is no net dipole strength, there will be some electrostatic interaction if the molecules are very close and are not spherically symmetric like noble gas atoms are.)
Chemguide [[1]] notes: “Surprisingly dipole-dipole attractions are fairly minor compared to dispersion [London] forces, and their effect can only really be seen if you compare two molecules with the same number of electrons and the same size.” One reason is that thermal motion tends to kill off the dipole attractions by messing up the alignment between molecules. But note that the dipole forces act on top of the London ones, so everything else being the same, the molecules with a dipole strength will be bound together more strongly.
When more than one molecular species is around, species with inherent dipoles can induce dipoles in other molecules that normally do not have them.
Another way molecules can be kept together in a solid is by what are called “hydrogen bonds.” In a sense, they too are dipole-dipole forces. In this case, the molecular dipole is created when the electrons are pulled away from hydrogen atoms. This leaves a partially uncovered nucleus, since an hydrogen atom does not have any other electrons to shield it. Since it allows neighboring molecules to get very close to a nucleus, hydrogen bonds can be strong. They remain a lot weaker than a typical chemical bond, though.
Key Points
- Even neutral molecules that do not want to create other bonds can be glued together by various “Van der Waals forces.”
- These forces are weak, though hydrogen bonds are much less so.
- The London type Van Der Waals forces affects all molecules, even noble gas atoms.
- London forces can be modeled using the Lennard-Jones potential.
- London forces are one of these weird quantum effects. Molecules with inherent dipole strength feature a more classically understandable version of such forces.