Wonders of Unstable Isotopes
Two atoms may have the same number of protons, and thus be of the same element, yet differ in their number of neutrons. Such atoms are called isotopes, atoms of the same element having different masses. Thus isotopes are atoms of the same element that have different masses due to differences in the number of neutrons they contain. The name comes from the Greek phrase ‘isos topos’, meaning “same place”: because they have the same atomic number, isotopes of the same element occupy the same position on the periodic table. Because the atomic number of any element is established, sometimes isotopes are represented simply with the mass number, thus: 93Ag. They may also be designated with a subscript notation indicating the number of neutrons, so that this information can be obtained at a glance without having to do the arithmetic. For the silver isotope shown here, this is written as Isotopes can also be indicated by simple nomenclature: for instance, carbon-12 or carbon-13.
Many isotopes are stable, meaning that they are not subject to radioactive decay, but many more are radioactive. The latter, also known as radioisotopes, play a significant role in modern life. Carbon-14, for instance, is used for estimating the age of objects within a relatively recent span of time—up to about 5,000 years—whereas geologists and other scientists use uranium-238 to date minerals of an age on a scale with that of the Earth. Concerns over nuclear power and nuclear weapons testing in the atmosphere have heightened awareness of the dangers posed by certain kinds of radioactive isotopes, which can indeed be hazardous to human life. However, the reality is that people are subjected to considerably more radiation from non-nuclear sources.
Radioactivity is a term used to describe a phenomenon whereby certain materials are subject to a form of decay brought about by the emission of high-energy particles or radiation. Forms of particles or energy emitted in radiation include alpha particles (positively charged helium nuclei); beta particles (either electrons or subatomic particles called positrons); or gamma rays, which occupy the highest energy level in the electromagnetic radiation emitted by the Sun. Radioactivity will be discussed below, but for the present, the principal concern is with radioactive properties as a distinguishing factor between the two varieties of isotope. Isotopes are either stable or unstable. The unstable variety, known as radioisotopes, is subject to radioactive decay, but in this context, “decay” does not mean what it usually does. A radioisotope does not “rot”; it decays by turning into another isotope of the same element or even into another element entirely. (For example, uranium-238 decays by emitting alpha particles, ultimately becoming lead-206.) A stable isotope, on the other hand, has already become what it is going to be, and will not experience further decay.
Most elements have between two and six stable isotopes. On the other hand, a few elements such as technetium have no stable isotopes. Twenty elements, inclusive of gold, fluorine, sodium, aluminum, and phosphorus, have only one stable isotope each. The element with the most stable isotopes is easy to remember because its name is almost the same as its number of stable isotopes such as tin, with 10. As for unstable isotopes, there are over 1,000, some of which exist in nature, but most of which have been created synthetically in laboratories. This number is not fixed; in any case, it is not necessarily important, because many of these highly radioactive isotopes last only for fractions of a second before decaying to form a stable isotope. Yet radioisotopes in general have so many uses, in comparison to stable isotopes, that they are often referred to simply as “isotopes.”
Before proceeding with a discussion of isotopes and their uses, it is necessary to address a point raised very often, when it is stated that some atoms do have the same numbers of neutrons and protons. In fact, nuclear stability is in part a function of neutron-to-proton ratio. Stable nuclei with low atomic numbers (up to about 20) have approximately the same number of neutrons and protons. For example, the most stable and abundant form of carbon is carbon-12, with six protons and six neutrons. Beyond atomic number 20 or so, however, the number of neutrons begins to grow: in other words, the lowest mass number is increasingly high in comparison to the atomic number. For example, uranium has an atomic number of 92, but the lowest mass number for a uranium isotope is not 184, or 92 multiplied by two; rather it is 218. The ratio of neutrons to protons necessary for a stable isotope creeps upward along the periodic table: for example, tin with an atomic number of 50, has a stable isotope with a mass number of 120, indicating a 1.4 to 1 ratio of neutrons to protons. For mercury-200, this ratio is 1.5 to 1.
By definition the higher is the atomic number the greater is the number of protons in the nucleus. This means that more neutrons are required to “bind” the nucleus together. In fact, all nuclei with 84 protons or more (i.e., starting at polonium and moving along the periodic table) are radioactive, for the simple reason that it is increasingly difficult for the neutrons to withstand the strain of keeping so many protons in place. One can predict the mode of radioactive decay by noting whether the nucleus is neutron-rich or neutron-poor. While neutron-rich nuclei undergo beta emission, which decreases the numbers of protons in the nucleus, neutron-poor nuclei typically undergo positron emission or electron capture, the first of these being more prevalent among the lighter nuclei. Elements with atomic numbers of 84 or greater generally undergo alpha emission, which decreases the numbers of protons and neutrons by two each.