Isotopes

Isotopes are variants of a particular chemical element. While all isotopes of a given element share the same number of protons, each isotope differs from the others in its number of neutrons. The term isotope is formed from the Greek roots isos (equal) and topos (place). Hence: "the same place," meaning that different isotopes of a single element occupy the same position on the periodic table. The number of protons within the atom's nucleus uniquely identifies an element, but a given element may in principle have any number of neutrons. The number of nucleons (protons and neutrons) in the nucleus, known as the mass number, is not the same for two isotopes of any element. For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of carbon is 6 which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7 and 8 respectively.

One of the most important scientific discoveries of the twentieth century was the discovery of isotopes: variants of elements that differ only in the number of neutrons contained in their nuclei, and therefore have different masses. This seemingly subtle variation has profound effects that can be exploited in numerous ways. Some isotopes may be unstable and can fission, leading to radioactivity and nuclear reactions. Isotopic substitution in specific locations in molecules can also have profound effects. Absorption and emission spectra may be altered: the wavelengths of features may shift. Reaction rates and equilibrium concentrations may change. Some properties may change dramatically: the effects of changing the number of particles in a nucleus can cause population of some atomic or molecular energy states to appear or disappear altogether. Nonzero nuclear spin in some isotopes (and not in others) open molecules to study by new methods, such as Nuclear Magnetic Resonance (NMR). All these effects can be exploited in order to obtain a window into detailed reaction mechanisms, reaction rates, and to effect chemical separations otherwise impossible.

The recognition of the great utility of isotopic substitution in the last three-quarters of a century has also spurred much theoretical work seeking to understand the effects of isotopic substitution and to predict new effects that can be exploited for chemical studies.



__ History: __

 * Isotopes were first suggested in 1912 based on radioactive decay chains
 * It was confirmed in 1913 by observing radioactive masses

__ Occurence in Nature: __

 * Elements are composed of one or more naturally occuring isotopes
 * Only 80 elements have any stable isotopes and 26 of these have only one stable isotope.
 * Thus, about 2/3 of stable elements occur naturally on Earth in multiple stable isotopes, with the largest numer of stable isotopes for an elesment being ten, for tin.

Isotopes are elements with the same number of protons, but varying numbers of neutrons.

Radioactive, primordial, and stable isotopes
Some isotopes are [|radioactive], and are therefore described as radioisotopes or [|radionuclides] , while others have never been observed to undergo radioactive decay and are described as [|stable isotopes]. For example, 14 C is a radioactive form of carbon while 12 C and 13 C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, [|[3]] of which 288 are [|primordial nuclides], meaning that they have existed since the solar system's formation. These include 33 nuclides with very long [|half-lives] (over 80 million years) and 255 which are formally considered as " [|stable isotopes] ", [|[3]] since they have not been observed to decay.

Many apparently "stable" isotopes are predicted by theory to be radioactive, with extremely long half-lives (this does not count the possibility of [|proton decay], which would make all nuclides ultimately unstable). Of the 255 nuclides never observed to decay, only 90 of these (all from the first 40 elements) are stable in theory to all known forms of decay. Element 41 ( [|niobium] ) is theoretically unstable via [|spontaneous fission], but this has never been detected. Many other stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay has yet been observed. The half-lives for these processes often exceed a million times the estimated age of the universe, and in fact there are 27 known radionuclides (see [|primordial nuclide] ) with half-lives longer than the age of the universe. Adding in the radioactive nuclides that have been created artificially, there are more than [|3100 currently known nuclides]. [|[4]] These include 905 nuclides which are either stable, or have half-lives longer than 60 minutes

Chemical and molecular 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 share a similar electronic structure. 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] ( 1   H ) and [|deuterium] (  2   H ), 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, which have more neutrons than lighter elements, the ratio of the nuclear mass to the collective electronic mass is far greater, and the relative mass difference between isotopes is much less. For these two reasons, the mass-difference effects on chemistry are usually negligible.

Isotope half-lives. Note that the plot for stable isotopes diverges from the line, protons Z = neutrons N as the element number Z becomes larger In similar manner, two [|molecules] that 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. As a consequence, 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.

Isotopes became known when scientists began researching radioactive decay chains which indicated more than 40 different species between Uranium and lead.