Isotope


Isotopes are distinct nuclear species of the same chemical element. They have the same atomic number and position in the periodic table, but different nucleon numbers due to different numbers of neutrons in their nuclei. While all isotopes of a given element have virtually the same chemical properties, they have different atomic masses and physical properties.
The term isotope comes from the Greek roots isos and topos, meaning "the same place": different isotopes of an element occupy the same place on the periodic table. It was coined by Scottish doctor and writer Margaret Todd in a 1913 suggestion to the British chemist Frederick Soddy, who popularized the term.
The number of protons within the atom's nucleus is called its atomic number and is equal to the number of electrons in the neutral atom. Each atomic number identifies a specific element, but not the isotope; an atom of a given element may have a wide range in its number of neutrons. The number of nucleons in the nucleus is the atom's mass number, and each isotope of a given element has a different mass number.
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.

Isotope vs. nuclide

A nuclide is a species of an atom with a specific number of protons and neutrons in the nucleus, for example, carbon-13 with 6 protons and 7 neutrons. Thus the terms are roughly synonymous, but the nuclide concept emphasizes nuclear properties over chemical properties, whereas the isotope concept emphasizes chemical over nuclear. The neutron number greatly affects nuclear properties, but its effect on chemical properties is negligible for most elements. Even for the lightest elements, whose ratio of neutron number to atomic number varies the most between isotopes, it usually has only a small effect although it matters in some circumstances. The term isotopes is intended to imply comparison. For example, the nuclides,, are isotopes, but,, are isobars. As the older and better-known term, isotope is however still used in some contexts where nuclide might be more appropriate, such as in nuclear technology and nuclear medicine.

Notation

An isotope/nuclide is specified by the name of the element followed by a hyphen and the mass number. When a chemical symbol is used, e.g. "C" for carbon, standard notation is to indicate the mass number with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left. Because the atomic number is already fixed by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript. The letter m is appended after the mass number to indicate a nuclear isomer, a metastable or energetically excited nuclear state, for example ; a number can be appended to it to distinguish different metastable states, though this is rare in practice.
The common pronunciation of the AZE notation is different from how it is written: is commonly pronounced helium-four instead of four-two-helium, and uranium two-thirty-five or uranium-two-three-five instead of 235-92-uranium or 235-uranium. This is not an error but the original spoken usage for isotope names, originating before AZE notation became established.

Radioactive, primordial, and stable isotopes

Some isotopes/nuclides are radioactive, and are therefore called radioisotopes or radionuclides, whereas others have never been observed to decay radioactively and are called stable isotopes or stable nuclides. For example, C is a radioactive form of carbon, while C and C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, of which 286 are primordial nuclides, meaning that they have existed since the Solar System's formation.
Primordial nuclides include 35 nuclides with very long half-lives and 251 that are considered "stable nuclides", as they have not been observed to decay. In most cases, if an element has stable isotopes, those isotopes predominate in the elemental abundance found on Earth and in the Solar System. However, in the cases of three elements the most abundant isotope found in nature is actually one extremely long-lived radioisotope of the element, despite these elements having one or more stable isotopes.
Theory predicts that many apparently "stable" nuclides are radioactive, with extremely long half-lives. Some stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay products have yet been observed, and so these isotopes are said to be "observationally stable". The predicted half-lives for these nuclides often greatly exceed the estimated age of the universe, and in fact, there are also 31 known radionuclides with half-lives longer than the age of the universe.
The total of all known nuclides, of which most have been created only artificially, is several thousand, of which 987 are stable or have a half-life longer than one hour; see List of nuclides.

History

Radioactive isotopes

The existence of isotopes was first suggested in 1913 by the radiochemist Frederick Soddy, based on studies of radioactive decay chains that indicated about 40 different species referred to as radioelements between uranium and lead, although the periodic table only allowed for 11 elements between lead and uranium inclusive.
Several attempts to separate these new radioelements chemically had failed. For example, Soddy had shown in 1910 that mesothorium, radium, and thorium X are impossible to separate. Attempts to place the radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive displacement law in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, whereas beta decay emission produced an element one place to the right. Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties.
Soddy proposed that several types of atoms could occupy the same place in the table. For example, the alpha-decay of uranium-235 forms thorium-231, whereas the beta decay of actinium-230 forms thorium-230. The term "isotope", Greek for "at the same place", was suggested to Soddy by Margaret Todd, a Scottish physician and family friend, during a conversation in which he explained his ideas to her. He received the 1921 Nobel Prize in Chemistry in part for his work on isotopes.
File:Discovery of neon isotopes.JPG|thumb|right|In the bottom right corner of J. J. Thomson's photographic plate are the separate impact marks for the two isotopes of neon: neon-20 and neon-22.
In 1914 T. W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to radiogenic variations in isotopic composition; the natural radioactive series ending with three different isotopes of lead.

Stable isotopes

The first evidence for multiple isotopes of a stable element was found by J. J. Thomson in 1912/1913 as part of his exploration into the composition of canal rays. Thomson channelled streams of ions through parallel magnetic and electric fields, measured their deflection by placing a photographic plate in their path, and computed their mass to charge ratio using a method that became known as the Thomson's parabola method. Each 'line' could be identified with a specific atomic weight, and therefore different elements and compounds could be identified. Thomson's initial paper in the Philosophical Magazine explains the technique used, but makes no comment on anomalous lines; however, in a talk given to the Royal Institution on 17 January 1913 Thomson identified "a line corresponding to an atomic weight 22, which can not be identified with the line due to any known gas". He went on to comment that
The origin of this line presents many points of interest; there are no known gaseous compounds of any of the recognized elements which have this molecular weight. Again, if we accept Mendeleef's Periodic Law, there is no room for a new element with this atomic weight.
The same lecture was then given to the Cambridge Philosophical Society on 27 January 1913, and later in the year Thomson presented his mature interpretation as the Royal Society's Bakerian Lecture. There he again showed the image of parabolic patches of light on the photographic plate, which suggested two species of neon nuclei with different mass-to-charge ratios. He wrote "There can, therefore, I think, be little doubt that what has been called neon is not a simple gas but a mixture of two gases, one of which has an atomic weight about 20 and the other about 22. The parabola due to the heavier gas is always much fainter than that due to the lighter, so that probably the heavier gas forms only a small percentage of the mixture."
F. W. Aston subsequently discovered multiple stable isotopes for numerous elements using a mass spectrograph, related to Thomson's method. In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to the integers 20 and 22, and that neither is equal to the known molar mass of neon gas. This is an example of Aston's whole number rule for isotopic masses, now known to be exceptionless, which states that large deviations of elemental molar masses from integers are due to the fact that the element is a mixture of isotopes. Aston similarly showed in 1920 that the molar mass of chlorine is a weighted average of the almost integral masses for the two isotopes 35Cl and 37Cl.