The atom was once considered indivisible, a concept rooted in the ancient Greek idea of atomos and solidified by John Dalton’s atomic theory in the early 1800s. For decades, the atom was imagined as a solid, featureless sphere, with its properties defined solely by its mass. This view held that the atom was the ultimate particle, incapable of being broken down into smaller components. This perception shattered at the turn of the 20th century with the discovery of subatomic particles—the electron, the proton, and the neutron—revealing the atom to be a complex structure of charged and neutral constituents.
The Discovery of the Electron
The first evidence that atoms were not indivisible came from experiments utilizing cathode ray tubes in the late 19th century. English physicist J.J. Thomson conducted a series of investigations using these evacuated glass tubes that contained two electrodes. When a high voltage was applied, a mysterious glow, the cathode ray, traveled from the negative cathode to the positive anode.
Thomson demonstrated that these rays were streams of tiny, negatively charged particles, not a form of light or energy. He showed this by applying both electric and magnetic fields, observing that the beam deflected toward the positive plate. This behavior confirmed the rays were composed of charged matter.
By balancing the forces of the electric and magnetic fields, Thomson was able to calculate the particle’s charge-to-mass ratio (\(e/m\)). The calculated ratio was startlingly large, indicating the particles were incredibly light. He determined that these particles, which he called “corpuscles,” were approximately 1,800 times lighter than the lightest known atom, hydrogen.
Crucially, the properties of the cathode rays remained the same regardless of the gas used inside the tube or the metal composing the electrodes. This universality proved that these negatively charged particles, later named electrons, were a fundamental component of all matter. To account for the atom’s overall electrical neutrality, Thomson proposed the “Plum Pudding” model, suggesting negative electrons were embedded in a diffuse sphere of positive charge.
Identifying the Positive Core
The Plum Pudding model was soon challenged by the experiments of Ernest Rutherford and his students, Hans Geiger and Ernest Marsden, beginning in 1909. Their famous Geiger-Marsden experiment, often called the gold foil experiment, sought to probe the internal structure of the atom using alpha particles. Alpha particles, known to be positively charged helium nuclei, were fired at an extremely thin sheet of gold foil.
Based on the prevailing model, the expectation was that the massive, positively charged alpha particles would pass straight through the uniform positive “pudding” with only minor deflections. The positive charge was thought to be too spread out to significantly alter the path of the fast-moving alpha particles. However, the results were entirely unexpected.
While the vast majority of alpha particles passed through the gold foil with little or no deflection, a small fraction—about one in every 8,000 particles—was deflected at large angles, some even bouncing directly backward. The result indicated that the positive charge and nearly all the atom’s mass were concentrated in an extraordinarily small, dense region.
Rutherford concluded that the atom must be mostly empty space, with its positive charge confined to a minute central core, which he named the nucleus. This dense, positively charged nucleus was responsible for the strong electrostatic repulsion that caused the large deflections. The positive charge within the nucleus was later identified as being carried by discrete particles, which Rutherford named protons in 1920, establishing the nuclear model of the atom.
The Search for the Neutral Component
Even after the discovery of the proton, a significant problem remained with the atomic model: the “missing mass” paradox. The atomic mass of every element heavier than hydrogen was roughly double the mass accounted for by its protons alone. For example, the helium nucleus contained two protons, but its atomic mass was four times that of a single proton. This suggested the nucleus contained other components.
Scientists, including Rutherford, speculated about the existence of a neutral particle with a mass similar to the proton. This particle was necessary to account for the extra weight without adding any charge, and to explain the existence of isotopes, which had the same chemical properties but different masses. The breakthrough came in 1932.
Earlier experiments by Walther Bothe and Herbert Becker, and later Irène and Frédéric Joliot-Curie, showed that when alpha particles were used to bombard beryllium, a highly penetrating, uncharged radiation was emitted. The Joliot-Curies observed that this radiation could eject high-energy protons from a block of paraffin wax, but they mistakenly interpreted this effect as powerful gamma rays.
James Chadwick repeated and refined these experiments, bombarding beryllium with alpha particles to produce the unknown radiation. He directed this radiation at various targets and analyzed the kinetic energy of the ejected protons. Chadwick demonstrated through the application of the conservation of energy and momentum that the observed energy transfer was physically impossible if the radiation consisted of massless gamma rays. The only explanation was that the beryllium emitted a neutral particle with a mass almost exactly equal to that of the proton. Chadwick had discovered the neutron, completing the picture of the atom’s basic structure and resolving the mystery of the missing atomic mass.