Rutherford's Gold-Foil Experiment: The Nuclear Model of the Atom

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.

The experiment was done to validate the prevailing atomic model then, the plum-pudding model (some refer to it as the raisin-bread model) championed by J.J. Thomson. According to Rutherford, if Thomson's atomic model is correct, each high-velocity alpha particle will just pass straight through the thin gold foil suffering, at most, minor deflections because of the weak influence of the spread-out ball of positive charge and the effect of the electrons being negligible.

What Geiger and Marsden actually observed was stunning. Most of the alpha particles did indeed traveled straight through the foil with little or no deviation. But a small fraction (about 1 in 10,000) rebounded, ending up on the same side of the foil as the incoming beam. A few were returned almost along the same tracks as they went in. Rutherford upon hearing of these rebounds, described it as the most incredible event of his life. It was, he said, "as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." Such huge deflections could mean only one thing: some of the alpha particles had run into massive concentrations of positive charge and, since like charges repel, had been hurled straight back by them. Thomson's plum-pudding model of the atom failed to explain these observations.

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.

Rutherford's nuclear model of the atom was considered as one giant step forward in understanding nature at the microscopic level. But even as it closed the issue on the alpha particle experiment, it threw open another one. Since the nucleus and its retinue of electrons are oppositely charged, and therefore attract one another, there didn't seem anything to stop the electrons from being pulled immediately into the nucleus. Rutherford countered by saying that the atom was like a miniature solar system: the electrons circled the nucleus in wide orbits just as planets orbit the sun. This is the picture of atoms that most of us still carry around in our heads. It's an appealing, easy-to-grasp image – one that's inspired many a logo of the atomic age. Yet theorists were well aware of its shortcomings right from the start.

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.