Methane has no such charges, and hence has little tendency for two molecules to stick together. The only attractions between molecules are the weak van der Waals forces mentioned previously for H₂ and He. These attractions arise because, although an atom in a molecule may be electrically nonpolar on the average over a finite period of time, at any given instant the electrons may not be distributed symmetrically around the nucleus. This is illustrated for three atoms at the right. The first drawing shows the time average, with a symmetrical distribution of electrons around each nucleus. The following three drawings show "snapshots" of the atoms at three instants in time when the random motion of electrons has brought about shortlived attractions between atoms A and B, B and C, and A and C.These attractions may seem small, but they are not insignificant. They are the forces between atoms in neighboring methane molecules that make methane gas finally condense to a liquid at -164C. The strengths of van der Waals forces depend mainly on surface areas of molecules. Hence gaseous H₂ molecules, which are smaller than CH₄ molecules, must be cooled to -253C before they move slowly enough that van der Waals attractions can make them stick to one another in a liquid.