London dispersion forces
/ LON-dun dis-PER-shun /
Even a perfectly symmetric molecule, with its electrons evenly spread, is not perfectly even from one instant to the next: the electron cloud jiggles, and for a fleeting moment more of it bunches on one side. That tiny, momentary lopsidedness tugs on the next molecule's electrons, and the two flicker in sympathy — a kind of synchronized shimmer that gently pulls them together. That pull is the London dispersion force.
London dispersion forces are the attraction that arises between any two molecules from these constantly forming, instantaneous dipoles and the matching dipoles they induce in their neighbors. They are present in every substance without exception, even noble gases and oils. Their strength grows with the number of electrons and the size of the molecule — bigger, floppier electron clouds are easier to distort, so heavier molecules attract more strongly.
Why they matter: dispersion forces are usually the largest part of the van der Waals attraction and are the only intermolecular force in nonpolar substances, so they alone explain why iodine is a solid, bromine a liquid, and chlorine a gas at room temperature. The honest caveat is that any single dispersion contact is extraordinarily weak; their dominance in large molecules comes from sheer numbers stacking up.
Going down the halogens, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid at room temperature — purely because larger atoms have more electrons and stronger London dispersion forces.
More electrons mean stronger dispersion forces and higher boiling points.
London dispersion is the one intermolecular force present in absolutely everything — even atoms and molecules with no permanent dipole. It is named after physicist Fritz London, who explained it using quantum mechanics; the word 'dispersion' refers to that theory, not to anything spreading out.