Isotope

Isotopes are atoms of a chemical element whose nuclei have the same atomic number, Z, but different atomic weights, A. The word isotope, meaning at the same place, comes from the fact that isotopes are located at the same place on the periodic table.

The atomic number corresponds to the number of protons in an atom. Thus, isotopes of a particular element contain the same number of protons. The difference in atomic weights results from differences in the number of neutrons in the atomic nuclei. In scientific nomenclature, isotopes are specified by the name of the particular element by a hyphen and the number of nucleons (protons and neutrons) in the atomic nucleus (e.g., helium-3, carbon-12, carbon-14, iron-57, uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g., 3He, 12C, 14C, 57Fe, 238U).

Collectively, the isotopes of the elements form the set of nuclides. A nuclide is a particular type of nucleus (characterised by A and Z). Strictly speaking, we should say that an element such as fluorine consists of one nuclide rather than that it has one isotope. Similarly, the tables at the foot of this article are tables of nuclides.

In a neutral atom, the number of electrons equals the number of protons. Thus, isotopes of a given element also have the same number of electrons and the same electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The primary exception is that, due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This "mass effect" is most pronounced for protium (1H) and deuterium (the common name of 2H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion and stabilizing the nucleus. For this reason neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus, for example, although the neutron/proton ratio of 3He is 1/2, the neutron/proton ratio of 238U is &gt;3/2. However, if too many neutrons are present, the nucleus becomes unstable.

Because isotopes of a given element have different numbers of neutrons they also have different neutron/proton ratios. This affects the nuclear stability, with the result that some isotopes are subject to nuclear decay. The decay of these radioactive isotopes (radioisotopes for short) is an important topic in nuclear physics. By studying the manner in which this decay occurs, physicists gain insight into the properties of the atomic nucleus.

In general, several isotopes of each element can be found in nature. Stable isotopes are by far the most abundant; however, significant quantities of long-lived unstable isotopes, such as uranium-238, can also be found. Small amounts of short-lived radioactive isotopes are also present in nature. These arise as products of the decay of larger long-lived radioactive nuclei. The atomic mass for an element in the periodic table is the average of the natural abundance of the isotopes of that element.

The amounts of the various isotopes on earth is ultimately the result of the amounts formed in stars and supernovae, and the subsequent decay patterns of the radioactive nuclei formed in these processes. After that, the formation of the solar system also influenced heavily on the proportions of different isotopes found here, since lighter nuclei was more easily blown away towards the outer parts of the solar system, by the solar wind immediately after the sun was formed. This is also why the gas giants are located further from the sun.

Applications of isotopes
Several applications exist that capitalize on properties of the various isotopes of a given element. One of the most common applications is as a tracer or marker in a technique called isotopic labeling. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectroscopy or infrared spectroscopy, which is mass sensitive because heavier atoms vibrate at different frequencies than lighter atoms.

An example of the use of isotopic labeling is the study of phenol (C6H5OH) in water. Upon adding phenol to deuterated water (water containing D2O in addition to the usual H2O), researchers observed the substitution of deuterium for the hydrogen in the hydroxyl group (C6H5OD), indicating that phenol readily undergoes hydrogen-exchange reactions with water. Only the hydroxyl group was affected, indicating that the other 5 hydrogen atoms did not participate in these exchange reactions.

Isotopic substitution can also be used to determine the mechanism of a reaction via the kinetic isotope effect.

In addition to isotopic labeling, several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D, 13C, and 31P. Mossbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as 57Fe.

Radioactive isotopes also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. The process of isotope separation represents a significant technological challenge.

Radioisotopes are also frequently used in medicine, biochemistry, and chemistry as tracers. Small quantities of the radioisotopes can be readily detected due to characteristic emissions by the decaying nuclei.

The natural radioactive decay of 14C enables Radiocarbon dating.

The relative stability of the 12C isotope means that it can be used be used in the definition of the Mole, an SI base unit.