Isotope

An isotope is one of several kinds of atoms of the same element that have different masses. These atoms have the same number of protons in their nuclei, but different numbers of neutrons, and therefore different mass numbers. The term isotope comes from the Greek isos topos, which means same place, because isotopes of the same element have the same atomic number and therefore occupy the same place in the periodic table.

British chemist Frederick Soddy (1877–1956) discovered and named isotopes in 1911 while he was working with New Zealand/British physicist Ernest Rutherford (1871–1937), although the neutron had not been identified yet. Soddy considered the difference to be one of radioactive properties, not atomic mass. He was awarded the 1921 Nobel Prize in chemistry for his work on the nature and occurrence of isotopes. In 1913, English physicist J. J. Thomson (1856–1940) found the first stable isotopes due to atomic mass difference while working with neon. English chemist Francis W. Aston (1877–1945) developed a mass spectrograph that he used to identify isotopes. He found 212 of the 281 naturally occurring stable isotopes and, thereafter, received the 1922 Nobel Prize in chemistry for this work.

Most elements have two or more isotopes, although there are 20 elements that have only one. Tin is the element with the largest number of isotopes; it occurs in nature as a mixture of 10 different isotopes ranging in mass from 112 to 124. A survey of the periodic table reveals that elements of even atomic numbers have more isotopes than do those of odd atomic number. Whether the nucleus has an odd or even number of neutrons seems to also enter into an atom’s ability to form isotopes. The greatest number of stable nuclei has an even number of both protons and neutrons; there are not as many stable nuclei with an even number of protons and an odd number of neutrons or an even number of neutrons and an odd number of protons. There are only four isotopes with odd numbers of both protons and neutrons. The shell model of the nucleus formulated by German-born American physicist Maria Goeppert-Mayer (1906–1972), for which she received the 1963 Nobel Prize in physics, was an attempt to explain the high number of isotopes with 2, 8, 20, 28, 50, and 82 protons or neutrons. Nuclei with these magic numbers of protons or neutrons have many stable isotopes.

The element carbon, for example, has two stable isotopes, carbon-12 and carbon-13, symbolized as ¹²C and ¹³C, respectively. The numbers 12 and 13 are the mass numbers of the isotopes—the total numbers of the protons plus neutrons in their nuclei. Because all carbon atoms have six protons in their nuclei, ¹³C must have seven neutrons (13-6) in its nucleus, and ¹²C has six (12-6). The element carbon as found in nature consists of 98.89% ¹²C atoms and 1.11% ¹¹C atoms. Carbon in living organisms contains also a very small amount of ¹⁴C, a radioactive isotope that is used in radiocarbon dating.

Most elements have between two and six stable isotopes (as opposed to unstable, or radioactive ones). Twenty elements, including fluorine, sodium, aluminum, phosphorus, and gold consist of only one stable isotope each. Tin, however, has ten—more than any other element. The number of stable isotopes an element has is determined by the relative stabilities of various numbers of neutrons and protons in their nuclei.

Only two isotopes have been given distinctive names, both isotopes of hydrogen. The stable isotope 2^H is known as deuterium, or heavy hydrogen, and the radioactive isotope 3^H is called tritium.

Because isotopes of the same element have identical chemical properties, they cannot be separated by chemical methods, but only by methods that are based on their mass differences, such as mass spectrometry. One of the extraordinary accomplishments of the Manhattan Project, which created the atomic bomb during World War II (1939–1945), was the successful separation of large amounts of ²³⁵U, the highly fissionable isotope of uranium, from the much more abundant ²³⁸U by allowing a gaseous uranium compound to diffuse through porous barriers. Being heavier, the ²³⁸U-containing molecules move more slowly through the barriers.

Over one thousand radioisotopes—radioactive isotopes—either exist in nature or have been made artificially by bombarding stable isotopes in particle accelerators. They are useful in so many applications that the word isotope is commonly used to mean radioisotope, as if stable isotopes did not exist.

Following the identification of the neutron by English physicist James Chadwick (1891–1974) in 1932, scientists experimented with reactions produced by neutrons. When the nucleus of most elements absorbed an additional neutron, the element changed into a heavier isotope of the same element. Although some were stable, this method led to the production of radioactive isotopes that do not occur in nature. It is thought that many of the chemical elements in the solar system were formed when neutrons were captured in the interior of the earlier stars, forming radioactive isotopes of elements that then changed to different elements through radioactive decay, transforming them into atoms of heavier elements. Proof of this was obtained with data from NASA’s Hubble Space Telescope during the 1990s, when eight elements heavier than zinc were discovered in interstellar gas: arsenic, gallium, germanium, krypton, lead, selenium, thallium, and tin.

Radioisotopes have become increasingly more important in many fields in the past 60 years or so. Some radioisotopes are produced by natural processes. Chlorine-36 is produced from chlorine-35 by thermal neutron activation; beryllium-10 is formed from atmospheric oxygen by cosmic ray action; car-bon-14 is produced in the atmosphere when nitrogen-14 nuclei collide with high energy neutrons that are in cosmic radiation. The ratio of carbon-14 to carbon-12 taken in by living cells depends on the amount of cosmic radiation coming through the atmosphere. Scientists consider this ratio to have remained (fairly) constant over geologic time; although it has been shown to vary with the solar cycle as well as artificially made activities that put high energy particles into the atmosphere, such as the detonation of atomic weapons or the explosion of the Chernobyl (Ukraine) nuclear power plant. Once living organisms die, they no longer take in carbon-14, while the carbon-14 already part of the organism will decay according to the normal half-life decay period of the element. This process lets scientists approximate the age of organic materials, a method known as radiocarbon dating for which American physical chemist Willard F. Libby (1908–1980) received the 1960 Nobel Prize in chemistry. Ratios of naturally occurring radioactive minerals to their decay daughters can be used in determining the age of geological materials. To determine the age of the oldest materials requires isotopes with long half-lives; for example, rubidium-87 decays to strontium-87, samarium-147 decays to neodymium-143, potassium-40 decays to argon-40, rhenium-187 decays to osmium-187, and uranium isotopes decay to stable isotopes of lead.

Many radioisotopes are produced in nuclear fission reactors either as fission fragments or decay daughters of fission fragments. Rubidium-86, cesium-136, and molybdenum-99, whose decay product technetium-99 is widely used in medical applications, are uranium-235 fission fragments. Other radioisotopes are produced in fission reactors through bombardment by the neutrons that were released in the fission chain reaction. An example of this reaction is the production from cobalt-59 to cobalt-60, a good source of gamma radiation for many applications. Other radioisotopes are produced in machines where particles like protons, deuterons, and alpha particles, are accelerated to very high energies by an electric current as they move within a magnetic field and then are released to crash into a non-radioactive target, changing it into a radioactive isotope. In 1930, British physicist Sir John Cockcroft (1897–1967) and British physicist Ernest T.S. Walton (1903–1995) developed a machine that accelerated protons; they received the 1951 Nobel Prize in physics for transmutation of atomic nuclei by artificially accelerated atomic particles. In 1931, American nuclear physicist Ernest Lawrence (1901–1958) designed the first cyclotron, for which he received the 1939 Nobel Prize in physics. Lawrence used his cyclotron to accelerate deuterons that were then crashed into sodium, producing a ther-apeutically useful radioisotope of sodium. Newer radioisotopes produced this way for medical applications include thallium-201 and iodine-123.

Neutron activation analysis is a method of producing radioisotopes in very small samples of a material. Based on the properties of the radioisotopes produced, the elements in the original material can be verified. This technique was developed by Edward Sayre and Heather Lechtman at the Brookhaven National Laboratory in 1969, and, has been successfully used to authenticate art works.

Radionuclides are now being used in many fields—in research, industry, agriculture, and medicine, to name a few—because they can be identified in different locations by instruments that are sensitive to radiation. As early as 1911, Hungarian-Danish chemist Georg von Hevesey (1859–1906) suggested that radioactivity could be used to label substances to make them easily recognizable. In 1934, he used a radioisotope of phosphorus to study how human tissue absorbs phosphate. He was awarded the 1943 Nobel Prize in chemistry for his work on tracers, which let an investigator follow the movement of a particular element. In medicine, tracers are used to diagnose and even treat some medical problems. A radioisotope is substituted for the stable form of a chemical that is normally used by a specific organ. Examples of widely used tracers in medicine are: calcium-47 for studies of bone formation, chromium-51 for red blood cell studies, cobalt-58 for diagnosing pernicious anemia, iodine-131 for the thyroid gland, chromium-51 for the spleen, gallium-67 for the lymph glands, phosphorus-32 for the liver, strontium-87 for bones, technetium-99 for the brain and liver, thal-lium7-201 for the heart, and xenon-133 for blood flow studies. American medical physicist Rosalyn Sussman Yalow (1921—) received the 1977 Nobel Prize in medicine and physiology for her work in developing radio-immunoassay procedures that use radionuclides in laboratory tests to diagnose medical problems.

See also Dating techniques; Mass number; Nuclear fission; Nuclear medicine; Proton; Radioactive tracers.

Isotope

Isotopes are two forms of an element with the same atomic number but different mass number. The existence of isotopes can be understood by reviewing the structure of atoms.

All atoms contain three kinds of basic particles: protons, neutrons, and electrons. (Hydrogen is the only exception to this statement; most hydrogen atoms contain no neutrons.) The protons and neutrons in an atom are found in the atomic nucleus, while the electrons are found in the space around the nucleus.

The number of protons in a nucleus defines an atom. Hydrogen atoms all have one proton in their nucleus; helium atoms all have two protons in their nucleus; lithium atoms all have three protons in their nucleus; and so on. The number of protons in an atom's nucleus is called its atomic number. Hydrogen has an atomic number of 1; helium, an atomic number of 2; and lithium, an atomic number of 3.

But atoms of the same element can have different numbers of neutrons. Some helium nuclei, for example, have two neutrons; others have only one. The mass number of an atom is the total number of protons and neutrons in the atom's nucleus. The two-neutron atom of helium has a mass number of four (two protons plus two neutrons). The one-neutron atom of helium has a mass number of three (two protons plus one neutron).

Another way of defining isotopes, then, is to say that they are different forms of an atom with the same number of protons but different numbers of neutrons.

Most elements have at least two stable isotopes. The term stable here means not radioactive. Twenty elements, including fluorine, sodium, aluminum, phosphorus, and gold, have only one stable isotope. By contrast, tin has the largest number of stable isotopes of any element, ten.

Representing isotopes

Isotopes are commonly represented in one of two ways. First, they may be designated by writing the name of the element followed by the mass number of the isotope. The two forms of helium are called helium-4 and helium-3. Second, isotopes may be designated by the chemical symbol of the element with a superscript that shows their mass number. The designations for the two isotopes of helium are ⁴He and ³He.

Radioactive isotopes

A radioactive isotope is an isotope that spontaneously breaks apart, changing into some other isotope. As an example, potassium has a radioactive isotope with mass number 40, ⁴⁰K or potassium-40. This isotope breaks down into a stable isotope of potassium, ³⁹K or potassium-39.

Radioactive isotopes are much more common than are stable isotopes. At least 1,000 radioactive isotopes occur in nature or have been produced synthetically in particle accelerators (atom-smashers) or nuclear reactors (devices used to control the release of energy from nuclear reactions).

Applications

Isotopes have many important applications in theoretical and practical research. The advantage of using two or more isotopes of the same element is that the isotopes will all have the same chemical properties but may differ from each other because of their mass differences. This difference allows scientists to separate one isotope from another. An important example of this process is the way isotopes were used to purify uranium during World War II (1939–45).

Two common isotopes of uranium exist, ²³⁵U and ²³⁸U. Of these two, ²³⁸U is much more abundant, making up about 99.3 percent of the uranium found in nature. But only ²³⁵U can be used in making nuclear weapons and nuclear reactors. Since both ²³⁵U and ²³⁸U have the same chemical properties, how can the valuable ²³⁵U be separated from the more abundant, but valueless, ²³⁸U?

One answer to this question was to convert natural uranium to a gas and then allow the gas to diffuse (spread out) through a porous barrier. Researchers found that the ²³⁵U in the natural uranium was slightly less heavy than the ²³⁸U, so it diffused through the barrier slightly more quickly. But because the difference in mass between the two isotopes is not very great, the diffusion had to be repeated many times before the two isotopes could be separated very well. Eventually, however, enough ²³⁵U was collected by this process to make the world's first nuclear weapons.

[See also Atomic mass; Carbon family; Dating techniques; Nuclear fission; Nuclear medicine; Periodic table; Radioactive tracers; Radioactivity; Spectrometry ]

Isotope

An isotope is one of several kinds of atoms of the same element that have different masses. These atoms have the same number of protons in their nuclei, but different numbers of neutrons, and therefore different mass numbers. The term isotope comes from the Greek isos topos, which means same place, because isotopes of the same element have the same atomic number and therefore occupy the same place in the periodic table .

The element carbon , for example, has two stable isotopes, carbon-12 and carbon-13, symbolized as 12C and 13C. The numbers 12 and 13 are the mass numbers of the isotopes—the total numbers of the protons plus neutrons in their nuclei. Because all carbon atoms have six protons in their nuclei, 13C must have seven neutrons (13-6) in its nucleus, and 12C has six (12-6). The element carbon as we find it in nature consists of 98.89% 12C atoms and 1.11% 11C atoms. Carbon in living organisms contains also a very small amount of 14C, a radioactive isotope that is used in radiocarbon dating.

Most elements have between two and six stable isotopes (as opposed to unstable, or radioactive ones). Twenty elements, including fluorine, sodium , aluminum , phosphorus , and gold consist of only one stable isotope each. Tin, however, has ten—more than any other element. The number of stable isotopes an element has is determined by the relative stabilities of various numbers of neutrons and protons in their nuclei.

Only two isotopes have been given distinctive names, both isotopes of hydrogen . The stable isotope 2H is known as deuterium , or heavy hydrogen, and the radioactive isotope 3H is called tritium .

Because isotopes of the same element have identical chemical properties, they cannot be separated by chemical methods, but only by methods that are based on their mass differences, such as mass spectrometry . One of the extraordinary accomplishments of the Manhattan Project, which created the atomic bomb during World War II, was the successful separation of large amounts of235U, the highly fissionable isotope of uranium , from the much more abundant 238U by allowing a gaseous uranium compound to diffuse through porous barriers. Being heavier, the 238U-containing molecules move more slowly through the barriers.

Over 1,000 radioisotopes—radioactive isotopes—either exist in nature or have been made artificially by bombarding stable isotopes in particle accelerators . They are useful in so many applications that the word isotope is commonly used to mean radioisotope, as if stable isotopes did not exist.

See also Dating techniques; Mass number; Nuclear fission; Nuclear medicine; Proton; Radioactive tracers.

Isotope

Different forms of atoms of the same element. Atoms consist of a nucleus, containing positively-charged particles (protons) and neutral particles (neutrons), surrounded by negatively-charged particles (electrons). Isotopes of an element differ only in the number of neutrons in the nucleus and hence in atomic weight. The nuclei of some isotopes are unstable and undergo radioactive decay . An element can have several stable and radioactive isotopes, but most elements have only two or three isotopes that are of any importance. Also, for most elements the radioactive isotopes are only of concern in material exposed to certain types of radiation sources. Carbon has three important isotopes with atomic weights of 12, 13, and 14. C-12 is stable and represents 98.9% of natural carbon. C-13 is also stable and represents 1.1% of natural carbon. C-14 represents an insignificant fraction of naturally-occurring carbon, but it is radioactive and important because its radioactive decay is valuable in the dating of fossils and ancient artifacts. It is also useful in tracing the reactions of carbon compounds in research.

See also Nuclear fission; Nuclear power; Radioactivity; Radiocarbon dating

isotope One of two or more varieties of a chemical element whose atoms have a common number of protons and electrons (i.e. their atomic number is the same) but which vary in the number of neutrons in their nucleus (i.e. their atomic weight, signified by their mass number, is different). For example, hydrogen exists in the forms ¹₁H (one proton, no neutron), ²₁H (deuterium: one proton, one neutron), and ³₁H (tritium: one proton, 2 neutrons). Water in which ²₁H replaces the more common ¹₁H is known as ‘heavy water’. There are 300 naturally occurring isotopes, but only 92 naturally occurring elements, and in nature elements often occur as a mixture of isotopes, with one form being the most common. Isotopes may be produced by various nuclear reactions and the products are frequently radioactive. There are three different ways of specifying an isotope; for example ²³⁵U, U-235, and uranium 235 all indicate the isotope of uranium with a mass number of 235. See also ISOTOPIC DATING.