Van der Waals forces, or London dispersion forces, are caused by temporary dipoles created when electron locations are lopsided. The electrons are constantly orbiting the nucleus, and by chance they could end up close together. The uneven concentration of electrons could make one side of the atom more negatively charged than the other; this creates a temporary dipole. With more electrons in an atom, the shells are further away from the nucleus, and these forces become stronger.
Van der Waals forces explain how nitrogen can be liquefied (Figure 1). Nitrogen gas is diatomic and nonpolar; its chemical notation is N2. Since both nitrogen atoms in the molecule have the same electronegativity, there is no dipole, and the molecule is non-polar. If there are no dipoles, what would make the nitrogen atoms stick together to form a liquid? Van der Waals forces are the answer: they allow otherwise non-polar molecules to have attractive forces. However, they are by far the weakest forces that hold molecules together.
These intermolecular forces are also sometimes called "dipole-induced dipole" or "momentary dipole" forces. Not all molecules are polar, and yet we know that there are also intermolecular forces between non-polar molecules, such as carbon dioxide. In non-polar molecules the electronic charge is evenly distributed overall; however, it is possible that at a particular moment in time, the electrons might not be evenly distributed. That is to say, the molecule will have a temporary dipole. In other words, each end of the molecule has a slight charge, either positive or negative. When this happens, molecules that are next to each other attract each other very weakly. These London forces are found in the halogens (e.g., F2 and I2), the noble gases (e.g., Ne and Ar), and in other non-polar molecules, such as carbon dioxide and methane.