In 1909, Hans Geiger and Ernest Marsden, two researchers in Ernest Rutherford's laboratory at the University of Manchester, fired a beam of alpha particles at a thin gold metal foil. An alphaparticle, identified and named a decade earlier by Rutherford, is one of the types of radiation given off by radioactive elements such as uranium. Because these particles are fast-moving and positively charged (they're now known to be high-speed Helium nuclei), Rutherford reasoned they will serve as a good probe of the atomic structure of matter.
After months of studying, in 1911, Rutherford, presented a new model of the atom in which all of the positive charge is crammed inside a tiny, massive nucleus about ten thousand times smaller than the atom as a whole. That's equivalent to a marble in the middle of a football stadium. The much lighter electrons, he assumed, are positioned outside the nucleus. Everyone was amazed at the implication of Rutherford’s proposal---the atoms of which planets, people, objects, and everything else are made consisted almost entirely of empty space.
The fatal flaw in Rutherford's model is that it contains charges that are accelerating. The charges are on the electrons and the acceleration is due to the electrons always changing direction as they move around their orbits. (Things accelerate when they change speed and/or direction.) Since Maxwell's time, scientists had known that accelerating charges radiate energy. What was to stop the orbiting electrons in Rutherford's atom quickly (in fact, in about one hundred-millionth of a second) losing all their energy and spiraling into the nucleus?
The answer came from a young Dane, Niels Bohr, who joined the team at Manchester for a six-month spell in 1912, shortly after Rutherford went public with his new vision of the atom. Bohr played a hunch. He knew about Planck's quantum. He knew there was no way to save an electron inside an atom from plummeting into the nucleus if it could give off energy continuously. And so he said simply that electrons inside atoms can’t radiate continuously. They can only radiate in lumps, and these lumps are the same as Planck's quanta. For a given type of atom, say hydrogen, there's a limited number of stable orbits that an electron can occupy. Each of these orbits corresponds to a whole multiple of the basic quantum. As long as an electron is in one of these orbits, its energy, contrary to whatever classical physics might say, stays the same. If it jumps from an outer (higher energy) orbit to an inner (lower energy) orbit, the energy difference between the two is given off as a quantum of light. Once the electron reaches the lowest energy orbit, it can't fall any further and is safe from the clutches of the nucleus.
