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 atoms of an element can have two or more different structures, which have the same atomic number (number of protons) but different mass numbers (number of protons plus neutrons). Based on these differences, the element can have different forms known as isotopes, each of which is made up of atoms with the same atomic structure. Isotopes that are radioactive are called radioisotopes.
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. The isotopes of a given element have nearly identical chemical properties but their physical properties show somewhat greater variation. Thus the process of isotope separation represents a significant technological challenge.
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.

Contents

1 Scientific nomenclature
2 Isotones, Isobars, Nuclear isomers
3 Variation in properties of isotopes
4 Occurrence in nature
5 Molecular mass of isotopes
6 Applications of isotopes

6.1 Use of chemical properties
6.2 Use of nuclear properties


7 External links
8 Credits



The properties of isotopes can be used for a variety of applications. Many people are aware that specific radioactive isotopes are used to produce nuclear power and nuclear weapons. In addition, radioactive isotopes or isotopes of different masses can be used as tracers in chemical and biochemical reactions, or to date geological samples. Also, several forms of spectroscopy rely on the unique nuclear properties of specific isotopes.

Scientific nomenclature
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. For example, carbon-12 and carbon-14 are isotopes of carbon; uranium-235 and uranium-238 are isotopes of uranium. Alternatively, the number of nucleons (protons and neutrons) per atomic nucleus may be denoted as a superscripted prefix attached to the chemical symbol of the element. Thus, the above examples would be denoted as 12C, 14C, 235U, and 238U, respectively.

Isotones, Isobars, Nuclear isomers
Isotopes are nuclides having the same atomic number (number of protons). They should be distinguished from isotones, isobars, and nuclear isomers.

Isotones are nuclides that have the same number of neutrons. For example, boron-12 and carbon-13 are isotones, because there are seven neutrons in each of their atoms.
Isobars are nuclides that have the same mass number (sum of protons plus neutrons). For example, carbon-12 and boron-12 are isobars. (In meteorology, however, an isobar is a line of constant pressure on a graph.)
Nuclear isomers are different excited states (energy states) of the same type of nucleus. A transition from one nuclear isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. (Nuclear isomers should not be confused with chemical isomers.)
Variation in properties of 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 atomic mass (Mr) of an element is determined by its nucleons. For example, carbon-12 has six protons and six neutrons, while carbon-14 has six protons and eight neutrons. When a sample contains two isotopes of an element, the atomic mass of the element is calculated by the following equation:




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


Here, 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 capitalize on properties of the various isotopes of a given element.

Use of chemical properties
One of the most common applications is known as "isotopic labeling"—the use of unusual isotopes as tracers or markers in chemical and biochemical reactions. For example, isotopes of different masses can be distinguished by techniques such as mass spectrometry or infrared spectroscopy (see "Properties"). Alternatively, if a radioactive isotope is used, it can be detected by the radiation it emits—a technique called radioisotopic labeling.
A technique similar to radioisotopic labeling is radiometric dating. Using the known half-life of an unstable element, one can estimate the amount of time that has elapsed since a known level of isotope came into existence. The most widely known example is radiocarbon dating, which is used to determine the age of carbon-containing materials.
The kinetic isotope effect can be used to determine the mechanism of a reaction, by substituting one isotope for another.
Use of nuclear properties
The nuclear reactions of certain radioactive isotopes are utilized for the production of nuclear power and nuclear weapons.
Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used 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.
External links
All links retrieved March 9, 2018.

Atomic weights of all isotopes

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