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Isotope
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Isotopes (Greek isos = "equal", tópos = "site, place") are any of the different types of atoms (nuclides) of the same chemical element, each having a different atomic mass (mass number). Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons. Therefore, isotopes have different mass numbers, which give the total number of nucleons, the number of protons plus neutrons.
A nuclide is any particular atomic nucleus with a specific number of an atom
Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons.
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Isotopes (Greek isos = "equal", tópos = "site, place") are any of the different types of atoms (nuclides) of the same chemical element, each having a different atomic mass (mass number). Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons. Therefore, isotopes have different mass numbers, which give the total number of nucleons, the number of protons plus neutrons.
A nuclide is any particular atomic nucleus with a specific number of an atom
Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons. Collectively, all the isotopes of all the elements form the set of nuclides. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. Isotope is better 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. For example, it is more correct to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope.
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 (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g. 3He, 12C, 13C, 131I and 238U).
About 339 nuclides occur naturally on Earth, of which 256 (about 75%) are stable (or, to be careful, have never been observed to decay; this note is necessary because many "stable" isotopes are predicted to be radioactive with very long half-lives). Counting the radioactive nuclides not found in nature that have been created artificially, more than 3100 nuclides are currently known.
History of the term
The term isotope was coined in 1913 by Margaret Todd, a Scottish doctor, during a conversation with Frederick Soddy (to whom she was distantly related by marriage). Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if several elements occupied each position in the periodic table. Hence Todd suggested the Greek term for "at the same place" as a suitable name. Soddy adopted the term and went on to win the Nobel Prize for Chemistry in 1921 for his work on radioactive substances.
Soddy's use of the word isotope was initially with regard to radioactive (unstable) atoms. However, in 1913, as part of his exploration into the composition of canal rays, J. J. Thomson channeled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. This was the first observation of different stable isotopes for an element. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.
Variation in properties between isotopes
Chemical and atomic properties
A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and the same electronic structure, and because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium (1H) vis-ŕ-vis deuterium (2H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements, the absolute mass of nucleus relative to electrons is far more, and the relative mass difference between isotopes is much less, and thus the mass-difference effects on chemistry are usually negligible.
Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.
Nuclear properties and stability
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. Neutrons also stabilize the nucleus because at short ranges they attract each other and protons equally by the strong nuclear force, and this extra binding force also offsets the electrical repulsion between protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, an increasing ratio of neutrons are needed to form a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 3He is 1:2, the neutron:proton ratio of 238U is greater than 3:2.
Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). Xenon is the only element which has nine stable isotopes. There is no element with exactly eight stable isotopes. See list of elements by nuclear stability for a complete list. Five elements have seven stable isotopes, eight have six stable isotopes, nine have five stable isotopes, nine have four stable isotopes, nine have three stable isotopes, 16 have two stable isotopes (counting Ta-180m as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic elements, having a single primordial stable isotope which dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well).. In total, there are 256 nuclides which have not been observed to decay (see List of elements by nuclear stability). For the 80 elements which have one or more stable isotopes, the average number of stable isotopes is 256/80 = 3.20 isotopes per element.
Other effects besides the bulk ratio of protons and neutrons affect nuclear stability. For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five nucleons from existing for long enough to serve as platforms for building up of heavier elements during fusion formation in stars (see triple alpha process). A similar pairing pattern shows in the fact that the 256 known stable nuclides contain only five that have both an odd number of protons and an odd number of neutrons (odd-odd nuclei): 2H, 6Li, 10B, 14N, 180mTa (the last one was predicted to decay but this process was never observed). Also, four long-lived radioactive odd-odd nuclides (40K, 50V, 138La, 176Lu) occur naturally. Most odd-odd nuclides are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.
Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture cross-sections and gamma spectroscopy and nuclear magnetic resonance properties.
Occurrence in nature
Elements are composed of one or more naturally occurring isotopes, which are normally stable. Some elements have unstable (radioactive) isotopes, either because their decay is so slow that a fraction still remains since they were created (examples: uranium, potassium), or because they are continually created through cosmic radiation (tritium, carbon-14) or by decay from an isotope in the first category (radium, radon).
As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two-thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (element number 50). There are about 94 elements found naturally on Earth (up to plutonium, element 94, inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimate that the elements which occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 256 of these naturally-occurring isotopes are stable in the sense of never having been observed to decay as of the present time. All the known stable isotopes occur naturally on Earth); the other 85 naturally-occurring isotopes are radioactive, but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. An additional ~ 2700 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.
The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. A good example is chlorine, having the composition 35Cl, 75.8%, and 37Cl, 24.2%, giving an atomic mass of 35.5. Values like this confounded scientists before the discovery of isotopes, as most light element atomic masses are close to integer multiples of hydrogen.
According to generally accepted cosmology only isotopes of hydrogen and helium, and traces of some isotopes of lithium, beryllium and boron were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously-produced isotopes.
The most common isotope of hydrogen has no neutrons at all.
(See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.
The atomic mass (Mr) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see Mass defect), to slightly different masses of neutron and proton, and to the mass of electron shell of the atom. One should take into account that the mass number is an integer dimensionless quantity, whereas the atomic mass is a real number expressed in atomic mass units. But the difference between these quantities is less then 1% in any case.
The atomic masses of naturally occurring isotopes of an element determine the atomic weight of the element. When the element contains N isotopes, the equation below is applied for the atomic weight M:
where M1, M2, ..., MN are the atomic masses of each individual isotope, and are the relative abundances of these isotopes.
Applications of isotopes
Several applications exist that capitalize on properties of the various isotopes of a given element.
Use of chemical and biological properties
- Isotope analysis is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration of food products. The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.
- Another common application 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"). For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called 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, but more so with heavy elements such as uranium or plutonium, than with lighter elements such as hydrogen, lithium, carbon, nitrogen, and oxygen. The lighter elements are commonly separated by gas diffusion of their compounds such as CO and NO. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectroscopy.
See also
- Atom
- Table of nuclides
- Isotopomer
- List of particles
- Isotopes are nuclides having the same number of protons; compare:
- Isotones are nuclides having the same number of neutrons.
- Isobars are nuclides having the same mass number, i.e. sum of protons plus neutrons.
- Nuclear isomers are different excited states of the same type of nucleus. A transition from one isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. (Not to be confused with chemical isomers.)
- Bainbridge mass spectrometer
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