Pieces of Atoms

The assumption that atoms are indivisible was challenged more than half a century after Dalton’s lifetime. In 1896, Sir Joseph John Thomson (J.J. Thomson) identified negatively charged particles that were much smaller and lighter than the smallest atoms. Thomson was experimenting with cathode rays—streams of negatively-charged particles (now known as electrons) that emanate from a high-voltage electrical terminal inside a vacuum tube. The stream of electrons would curve in the presence of a magnet, because charged particles experience a sideways force when moving through a magnetic field. (See this section of chapter 2.) The cathode rays would also curve, at a different angle, in the presence of an electric field. By measuring the deflection of the rays in a magnetic field, and comparing this to the angle of deflection in an electric field, Thomson was able to estimate the mass-to-charge ratio of the electrons. To his surprise, he found that the mass-to-charge ratio of an electron is more than 1,000 times smaller than that of a hydrogen ion (an electrically charged atom of hydrogen). Since hydrogen atoms are the smallest of all atoms, this implied that cathode rays do not consist of atoms, but of particles much smaller than the smallest atoms.

The fact that there are particles much smaller than atoms suggested that atoms are not the fundamental building blocks of matter, but are made up of even smaller components. If atoms contain electrons, which have negative charge, then there must also be a positively charged ingredient in each atom, since the total charge of an atom is neutral. Moreover, the remaining portion of the atom must be relatively massive, since even the smallest atoms have thousands of times more mass than electrons. Thomson hypothesized that an atom consists of a positively charged sphere with negatively charged electrons scattered through it like raisins in plum pudding. His hypothesis about the structure of atoms is known as the plum pudding model.

Shortly after Thomson’s discovery of electrons, his former student Ernest Rutherford discovered that some kinds of atoms spontaneously break apart in a process called radioactive decay. (We’ll learn about that process in the next chapter.) When atoms of uranium, radium, and numerous other elements decay, they emit positively charged particles, which Rutherford called alpha particles. By measuring the deflection of alpha particles in a magnetic field, Rutherford was able to estimate the mass of these particles, just as Thomson had done with electrons. He found that alpha particles are relatively massive, about four times heavier than hydrogen atoms (and thousands of times heavier than electrons). Rutherford later determined that alpha particles consist of two protons and two neutrons. We’ll learn more about that discovery on the next page.

Rutherford noticed that alpha particles could travel through solid materials, if the materials were thin enough. They could even pass through thin sheets of gold, a very dense metal with its atoms tightly packed together. Rutherford asked one of his colleagues, Hans Geiger, to help him investigate this phenomenon. They designed a series of experiments, the most famous of which was conducted in 1909 by Geiger and an undergraduate student named Ernest Marsden. In the Geiger-Marsden experiment (sometimes called the Rutherford gold foil experiment), alpha particles emitted from a radioactive substance were directed toward a sheet of gold foil. A fluorescent screen was used to detect the particles after they passed through the foil. Whenever an alpha particle struck the screen, it would produce a tiny flash of light. Though these flashes were not visible to the naked eye, they could be seen using a microscope in a dark room.

The microscope and fluorescent screen were rotated around the foil at various angles, to see how the foil affected the paths of the alpha particles. Most of the particles either passed straight through the foil or were deflected at very small angles, but occasionally one would fly off in a different direction. Once in a great while, an alpha particle would even bounce off of the foil rather than passing through it!

These observations could not be explained by Thomson’s plum pudding model. Since alpha particles usually passed through the tightly-packed atoms in the foil, they obviously weren’t bouncing off of the spheres that were supposed to comprise the bulk of each gold atom. The few particles that did bounce back must have hit something inside the gold atoms. But according to Thomson’s model, the only things inside an atom are electrons. Electrons could not have caused the alpha particles to bounce back, since electrons have negative charge and would attract rather than repel the positively-charged alpha particles. Moreover, even if there were some unknown force that caused electrons and alpha particles to repel each other at close range, the relatively massive alpha particles would have easily pushed the tiny electrons out of the way.

Rutherford, who had previously accepted Thomson’s plum pudding model, was astonished by the results of the Geiger-Marsden experiment:

It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you.Quoted in Laylin K. James (ed.), Nobel Laureates in Chemistry, 1901-1992 (Washington, D.C.: American Chemical Society, 1993), 57.

To make sense of these surprising observations, Rutherford suggested that most of an atom’s mass is concentrated in a tiny, positively charged nucleus at the center of the atom. Alpha particles were deflected at sharper angles when they passed near an atom’s nucleus and were repelled by its positive charge. The reason some particles bounced off of the foil was that they happened to collide with a nucleus head-on. This happened very rarely: Geiger and Marsden estimated that only 1 in 8000 alpha particles bounced off the foil; the rest would pass through.See page 499 of their original report.

Because collisions with the nucleus were so uncommon, Rutherford concluded that the nucleus is extremely tiny compared to the size of the atom as a whole. He was right about that! Later experiments determined that the diameter of a typical nucleus is about 100,000 times smaller than the diameter of the atom as a whole. If you made a scale model of an atom, with a golf ball representing the nucleus, the model would be roughly 2.6 miles across!