Isotope

For a given chemical element, every atom has the same number of protons in its nucleus, but the number of neutrons per atom may vary. In other words, the element can have several different forms, which have the same atomic number (number of protons) but different mass numbers (number of protons plus neutrons). These different forms of the element are known as isotopes. The term isotope comes from Greek and means "at the same place"—all the different isotopes of an element are placed at the same location on the periodic table.

A particular atomic nucleus with a specific number of protons and neutrons is called a nuclide. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. Isotope is usually used when referring to several different nuclides of the same element; nuclide is more generic and is used when referencing only one nucleus or several nuclei of different elements.

In scientific nomenclature, isotopes and nuclides are specified by the name of the particular element (implicitly giving the atomic number) followed by a hyphen and the mass number. Examples are carbon-12, carbon-14, uranium-235, and uranium-238. Alternatively, the number of nucleons (protons and neutrons) per atomic nucleus is denoted as a superscripted prefix to the chemical symbol of the element. Examples are ¹²C, ¹⁴C, ²³⁵U, and ²³⁸U.

Variation in properties between isotopes

A neutral atom has the same number of electrons as protons. Thus, the atoms of all the isotopes of an element have the same number of protons and electrons and the same electronic structure. Given that the chemical behavior of an atom is largely determined by its electronic structure, the isotopes of a particular element exhibit nearly identical chemical behavior. The main exception to this rule is what is called the "kinetic isotope effect": heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element.

This "mass effect" is most pronounced for protium (¹H) as compared with deuterium (²H), because deuterium has twice the mass of protium. For heavier elements, the differences between the atomic masses of the isotopes are not so pronounced, and the mass effect is much smaller, usually negligible.

Likewise, two molecules that differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structures. Therefore, their physical and chemical properties will be almost indistinguishable (again with deuterium being the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Given that vibrational modes allow a molecule to absorb photons of corresponding (infrared) energies, isotopologues have different optical properties in the infrared range.

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. As protons are positively charged, they repel one another. Neutrons, being electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion. Neutrons also stabilize the nucleus, because at short ranges they attract each other and protons equally by the strong nuclear force, and this attraction also offsets the electrical repulsion between protons. For this reason, one or more neutrons are necessary for two or more protons to be bound together in a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus. For example, the neutron/proton ratio of ³He is 1:2, but the neutron/proton ratio of ²³⁸U is greater than 3:2. If the atomic nucleus contains too many or too few neutrons, it is unstable and subject to nuclear decay.

Occurrence in nature

Most elements have several different isotopes that can be found in nature. The relative abundance of an isotope is strongly correlated with its tendency toward nuclear decay—short-lived nuclides decay quickly and their numbers are reduced just as fast, while their long-lived counterparts endure. This, however, does not mean that short-lived species disappear entirely—many are continually produced through the decay of longer-lived nuclides. Also, short-lived isotopes such as those of promethium have been detected in the spectra of stars, where they are presumably being made continuously, by a process called stellar nucleosynthesis. The tabulated atomic mass of an element is an average that takes into account the presence of multiple isotopes with different masses and in different proportions.

According to generally accepted cosmology, virtually all nuclides—other than isotopes of hydrogen and helium, and traces of some isotopes of lithium, beryllium, and boron—were built in stars and supernovae. Their respective abundances result from the quantities formed by these processes, their spread through the galaxy, and their rates of decay. After the initial coalescence of the solar system, isotopes were redistributed according to mass (see also Origin of the Solar System). The isotopic composition of elements is different on different planets, making it possible to determine the origin of meteorites.

Molecular mass of isotopes

The molecular mass (Mr) of an element is determined by its nucleons. For example, Carbon-12 has 6 Protons and 6 Neutrons. When a sample contains two isotopes the equation below is applied:

$Mr={\frac {Mr(1)*\%abundance+Mr(2)*\%abundance}{100}}$

Where Mr(1) and Mr(2) are the molecular masses of each individual isotope, and %abundance is the percentage abundance of that isotope in the sample.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element.

Use of chemical properties

One of the most common applications is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy (see "Properties"). If radioactive isotopes are used, they can be detected by the radiation they emit (this is radioisotopic labeling).

A technique similar to radioisotopic labelling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.

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

Use of nuclear properties

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,¹⁵N, ¹³C, and ³¹P.

Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as ⁵⁷Fe.

Radionuclides 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.

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