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Modeled on an analogy to a liquid drop, the first term represents the favourable contribution to the binding of the nucleus made by short-range, attractive nuclear forces between neutrons and protons.The second term corrects the first by allowing for the expectation that nucleons at the surface of the nucleus, unlike those in the interior, do not experience forces of nuclear attraction equally from all sides.Generalizing from these and other data, English chemist Frederick Soddy in 1910 observed that “elements of different atomic weights [now called atomic masses] may possess identical (chemical) properties” and so belong in the same place in the periodic table.With considerable prescience, he extended the scope of his conclusion to include not only radioactive species but stable elements as well.A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. Uranium ores, for example, yielded ionium, and thorium ores gave mesothorium.
He expected a difference because uranium and thorium decay into different isotopes of lead.His work grew out of the study of positive rays (sometimes called canal rays), discovered in 1886 by Eugen Goldstein and soon thereafter recognized as beams of positive ions. The ions in the heavier ray had masses about two units, or 10 percent, greater than the ions in the lighter ray.To prove that the lighter neon had a mass very close to 20 and that the heavier ray was indeed neon and not a spurious signal of some kind, Aston had to construct an instrument that was considerably more precise than any other of the time.Isotopes are said to be stable if, when left alone, they show no perceptible tendency to change spontaneously.Under the proper conditions, however, say in a nuclear reactor or particle accelerator or in the interior of a star, even stable isotopes may be transformed, one into another.