The discovery of the neutron and its properties was central to the extraordinary developments in atomic physics in the first half of the 20th century. Early in the century, Ernest Rutherford used alpha particle scattering to discover that an atom has its mass and electric charge concentrated in a tiny nucleus. By 1920, isotopes of chemical elements had been discovered, the atomic masses had been determined to be approximately integer multiples of the mass of the hydrogen atom, and the atomic number had been identified as the charge on the nucleus.Byrne, J. Neutrons, Nuclei, and Matter, Dover Publications, Mineola, New York, 2011, Throughout the 1920s, the nucleus was viewed as composed of combinations of protons and electrons, the two elementary particles known at the time, but that model presented several experimental and theoretical contradictions.
The essential nature of the atomic nucleus was established with the discovery of the neutron by James Chadwick in 1932 and the determination that it was a new elementary particle, distinct from the proton.
The uncharged neutron was immediately exploited as a new means to probe nuclear structure, leading to such discoveries as the creation of new radioactive elements by neutron irradiation (1934) and the fission of uranium atoms by neutrons (1938). The discovery of fission led to the creation of both nuclear power and nuclear weapons by the end of World War II. Both the proton and the neutron were presumed to be elementary particles until the 1960s, when they were determined to be composite particles built from quarks.
Discovery of radioactivity At the start of the 20th century, the vigorous debate as to the existence of atoms had not yet been resolved. Philosophers such as Ernst Mach and Wilhelm Ostwald denied the existence of atoms, viewing them as a convenient mathematical construct, while scientists such as Arnold Sommerfeld and Ludwig Boltzmann saw that physical theories required the existence of atoms.
Radioactivity was discovered in 1896 by the French scientist Henri Becquerel, while working with phosphorescent materials. In 1898, Ernest Rutherford at Cavendish Laboratory distinguished two types of radioactivity, alpha rays and beta rays, which differed in their ability to penetrate, or travel into, ordinary objects or gases. Two years later, Paul Villard discovered gamma rays, which possessed even more penetrating power. These radiations were later identified with known particles: beta rays were shown to be electrons by Walter Kaufmann in 1902, alpha rays were shown to be helium ions by Rutherford and Thomas Royds in 1907, and gamma rays were shown to be electromagnetic radiation, that is, a form of light, by Rutherford and Edward Andrade in 1914. These radiations had also been identified as emanating from atoms, hence they provided clues to processes occurring within atoms. Conversely, the radiations were also recognized as tools that could be used in scattering experiments to probe the interior of atoms.
Gold foil experiment and the discovery of the atomic nucleus
At the University of Manchester between 1908 and 1913, Rutherford directed Hans Geiger and Ernest Marsden in a series of experiments to determine what occurs when alpha particles scatter from metal foil. Now called the Rutherford gold foil experiment, or the Geiger–Marsden experiment, these measurements made the extraordinary discovery that although most alpha particles passing through a thin gold foil experienced little deflection, a few scattered to a high angle. The scattering indicated that some of the alpha particles ricocheted back from a small, but dense, component inside the atoms. Based on these measurements, it was apparent to Rutherford by 1911 that the atom consisted of a small, massive nucleus with positive charge. The concentrated atomic mass was required to provide the observed deflection of the alpha particles, and Rutherford developed a mathematical model that accounted for the scattering.
While the Rutherford model was largely ignored at the time, when Niels Bohr joined Rutherford's group, he developed the Bohr model for electrons orbiting the nucleus in 1913. The Bohr model eventually led to an atomic model based on quantum mechanics by the mid-1920s.
Discovery of isotopes Concurrent with the work of Rutherford, Geiger, and Marsden, the radiochemist Frederick Soddy at the University of Glasgow was studying chemistry-related problems on radioactive materials. Soddy had worked with Rutherford on radioactivity at McGill University. By 1910, about 40 different radioactive elements, referred to as radioelements, had been identified between uranium and lead, although the periodic table only allowed for 11 elements. Every attempt to chemically isolate the radioelements mesothorium or thorium X from radium failed. Soddy concluded that these element were chemically the same element. At the suggestion of Margaret Todd, Soddy called these chemically identical elements isotopes. In 1913, Soddy and theorist Kazimierz Fajans independently found the displacement law, that an element undergoing alpha decay will produce an element two places to the left in the periodic system and an element undergoing beta decay will produce an element one place to the right in the periodic system. For his study of radioactivity and the discovery of isotopes, Soddy was awarded the 1921 Nobel Prize in Chemistry.
Prior to 1919 only atomic weights averaged over a very large number of atoms was available. In that year, Francis Aston built the first mass spectrograph, an improved form of a device built by J. J. Thomson to measure the deflection of positively charged atoms by electric and magnetic fields. Aston was then able to separate the isotopes of many light elements including neon, and . Aston discovered the isotopes matched William Prout's whole number rule: the mass of every isotope is a whole number multiple of hydrogen.Aston, Francis William. Mass spectra and isotopes. London: Edward Arnold, 1942.
Significantly, the one exception to this whole number rule was hydrogen itself, which had a mass value of 1.008. The excess mass was small, but well outside the limits of experimental uncertainty. Aston and others realized this difference was due to the binding energy of atoms. When a number of hydrogen atoms are bound into a atom, that atom's energy must be less than the sum of the energies of the separate hydrogen atoms. That lost energy, according to the mass-energy equivalence principle, means the atomic mass will be slightly less than the sum of the masses of its components. Aston's work on isotopes won him the 1922 Nobel Prize in Chemistry for the discovery of isotopes in a large number of non-radioactive elements, and for his enunciation of the whole number rule.
Atomic number and Moseley's law
Before 1913, chemists adhered to Mendeleev's principle that chemical properties derived from atomic weight. However, several places in the periodic table were inconsistent with this concept. For example cobalt and nickel seemed reversed. There were also attempts to understand the relationship between the atomic mass and nuclear charge. Rutherford knew from experiments in his lab that helium must have a nuclear charge of 2 and a mass of 4; this 1:2 ratio was expected to hold for all elements. In 1913 Antonius van den Broek hypothesized that the periodic table should be organized by charge, denoted by Z, not atomic mass and that Z was not exactly half of the atomic weight for elements. This solved the cobalt-nickel issue. Placing cobalt (Z=27, mass of 58.97), before the less heavy nickel (Z=28, mass of 58.68) gave the ordering expected by chemical behavior.
Henry Moseley set out to test Broek's hypothesis by measuring the electromagnetic emission spectra of heavier elements, such as cobalt and nickel, to see if they followed the ordering by weight or by atomic number. In 1913–1914 Moseley tested the question experimentally by using X-ray spectroscopy. He found that the most intense short-wavelength line in the X-ray spectrum of a particular element, known as the K-alpha line, was related to the element's position in the periodic table, that is, its atomic number, Z. Moseley found that the frequencies of the radiation were related in a simple way to the atomic number of the elements for a large number of elements.
Within a year, it was noted that the equation for the relation, now called Moseley's law, could be explained in terms of the 1913 Bohr model with reasonable extra assumptions about atomic structure in other elements. Moseley's result, by Bohr's later account, not only established atomic number as a measurable experimental quantity, but gave it a physical meaning as the positive charge on the atomic nucleus. The elements could be ordered in the periodic system in order of atomic number, rather than atomic weight. The result tied together the organization of the periodic table, the Bohr model for the atom, and Rutherford's model for alpha scattering from nuclei. It was cited by Rutherford, Bohr, and others as a critical advance in understanding the nature of the atomic nucleus.
Further research in atomic physics was interrupted by the onset of World War I. Moseley was killed in 1915 at the Battle of Gallipoli, while Rutherford's student James Chadwick was interned in Germany for the duration of the war, 1914–1918. In Berlin, Lise Meitner's and Otto Hahn's research work on determining the radioactive decay chains of radium and uranium by precise chemical separation was interrupted. Meitner spent much of the war working as a radiologist and medical X-ray technician near the Austrian front, while Hahn, a chemist, worked on research in poison gas warfare.
Rutherford's conjecture
In 1920, Rutherford gave a Bakerian lecture at the Royal Society entitled the "Nuclear Constitution of Atoms", a summary of recent experiments on atomic nuclei and conclusions as to the structure of atomic nuclei. By 1920, the existence of electrons within the atomic nucleus was widely assumed. It was assumed the nucleus consisted of hydrogen nuclei in number equal to the atomic mass number. But, since each hydrogen nucleus had charge +1 e, the nucleus required a smaller number of "internal electrons" each of charge −1 e to give the nucleus its correct total charge. The mass of protons is about 1800 times greater than that of electrons, so the mass of the electrons is incidental in this computation. Such a model was consistent with the scattering of alpha particles from heavy nuclei, as well as the charge and mass of the many isotopes that had been identified. There were other motivations for the proton-electron model. As noted by Rutherford at the time, "We have strong reason for believing that the nuclei of atoms contain electrons as well as positively charged bodies ...", namely, it was known that beta radiation was electrons emitted from the nucleus.