<P> In the Geiger--Marsden experiment, Hans Geiger and Ernest Marsden (colleagues of Rutherford working at his behest) shot alpha particles at thin sheets of metal and measured their deflection through the use of a fluorescent screen . Given the very small mass of the electrons, the high momentum of the alpha particles, and the low concentration of the positive charge of the plum pudding model, the experimenters expected all the alpha particles to pass through the metal foil without significant deflection . To their astonishment, a small fraction of the alpha particles experienced heavy deflection . Rutherford concluded that the positive charge of the atom must be concentrated in a very tiny volume to produce an electric field sufficiently intense to deflect the alpha particles so strongly . </P> <P> This led Rutherford to propose a planetary model in which a cloud of electrons surrounded a small, compact nucleus of positive charge . Only such a concentration of charge could produce the electric field strong enough to cause the heavy deflection . </P> <P> The planetary model of the atom had two significant shortcomings . The first is that, unlike planets orbiting a sun, electrons are charged particles . An accelerating electric charge is known to emit electromagnetic waves according to the Larmor formula in classical electromagnetism . An orbiting charge should steadily lose energy and spiral toward the nucleus, colliding with it in a small fraction of a second . The second problem was that the planetary model could not explain the highly peaked emission and absorption spectra of atoms that were observed . </P> <P> Quantum theory revolutionized physics at the beginning of the 20th century, when Max Planck and Albert Einstein postulated that light energy is emitted or absorbed in discrete amounts known as quanta (singular, quantum). In 1913, Niels Bohr incorporated this idea into his Bohr model of the atom, in which an electron could only orbit the nucleus in particular circular orbits with fixed angular momentum and energy, its distance from the nucleus (i.e., their radii) being proportional to its energy . Under this model an electron could not spiral into the nucleus because it could not lose energy in a continuous manner; instead, it could only make instantaneous "quantum leaps" between the fixed energy levels . When this occurred, light was emitted or absorbed at a frequency proportional to the change in energy (hence the absorption and emission of light in discrete spectra). </P>

Who was the first scientist to discover that atoms are divisible