Isotope Definition – Stable and Radioactive Isotopes

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What is an isotope? The symbol of isotopes, information on stable and radioactive isotopes, separation of isotopes and history.

Isotope Definition - Stable and Radioactive IsotopesIsotope; a particular class of one of the elements. All of the isotopes of an element have the same atomic number as the element, but they differ from each other in mass number. The atomic number is the number of protons in the nucleus of an, atom, ranging from 1 for hydrogen to 103 for lawrencium. The mass number is the total number of protons and neutrons in the nucleus of an atom. Because the atomic number also represents the number of electrons around the nucleus of a neutral atom, all isotopes of the same element have almost identical chemical properties.

Isotopes of the same element have the same number of protons in their nuclei, but they differ in the number of neutrons in their nuclei. For example, the three isotopes of hydrogen are H^1 (protium), H^2 (deuterium), and H^3 (tritium); the nucleus of protium has only one proton, the nucleus of deuterium has one proton and one neutron, and the nucleus of tritium has one proton and two neutrons. In the notation, the atomic number is written as a subscript, and the mass number is written as a superscript.

STABLE AND RADIOACTIVE ISOTOPES

It is estimated that the total number of isotopes now known is about 1,500; of these, about 280 are naturally occurring stable isotopes, about 25 are naturally occurring radioactive isotopes, and about 1,200 are artificially produced radioactive isotopes. Stable isotopes and radioactive isotopes are produced by separating them from a mixture of naturally occurring isotopes or by forming them in a nuclear reactor or with a particle accelerator. Tritium and plutonium-239, for example, are made as by-products in nuclear reactors.

Stable Isotopes: Some elements, among which are beryllium, sodium, and cobalt, have only one stable isotope. For the other elements, the number of stable isotopes ranges from two to ten; for example, chlorine has two, sulfur has four, calcium has six, and tin has ten. Tin has more stable isotopes than any other element.

The stable isotopes have several striking features. Most of the low-mass elements have only one or two stable isotopes. In these isotopes, the number of neutrons in each nucleus is almost the same as the number of protons. Many of the heavier elements have two to ten stable isotopes. In these isotopes, nuclear stability requires a larger ratio of neutrons to protons. For instance, the nucleus of one isotope of tin has 74 neutrons and 50 protons. In general, nuclear stability requires that the neutron-to-proton ratio increases as the atomic numbers of the elements increase.

Another striking feature is that an element with an odd atomic number has only one or two stable isotopes, whereas an, element with an even atomic number may have as many as ten isotopes. Apparently nuclear stability is enhanced by the pairing of protons; this is also true of the neutrons.

Radioactive Isotopes: Every element has at least one radioactive isotope. The radioactive isotope of hydrogen, for example, is tritium, which has a half-life of 12.5 years. Some of the heavier elements, such as platinum, have as many as 30 radioactive isotopes.

About 1,200 different radioactive isotopes have been produced by nuclear bombardment in nuclear reactors or particle accelerators. The further their neutron-to-proton ratio varies from
that of the stable isotopes, the shorter is their lifetime. Artificially produced radioactive isotopes have half-lives in the range from fractions of a second to years.

Some naturally occurring radioactive isotopes, such as some isotopes of the elements with atomic numbers from 84 to 92, have half-lives so long-over one billion years—that they have not completely decayed since they were formed. Systematic analysis of the abundance and half-lives of these isotopes provides information about the age of the earth and of the universe.

Separation of Isotopes. One method for the separation of isotopes is gaseous diffusion. Diffusion of a gaseous compound through a porous material results in a very slight concentration of the lighter isotope in a mixture, such as U^235 ( uranium-235 ) in a mixture of U^238 and U^235. The diffusion process can be cascaded so that the enriched fraction passes on to a succeeding stage, and the rest feeds back to earlier stages. This technique, using uranium hexafluoride as the gas, was used on a large scale during World War II to obtain greater concentrations of U^235 than the 0.7% concentration that occurs naturally.

Other methods of separation include differential electrolysis, evaporation, chemical exchange, and, electromagnetic separation. The last provides very pure samples of a particular isotope, but only in very small amounts.

Uses: Separated isotopes or enriched mixtures have been widely used in industry and research. Among these, uranium-235 is useful because it undergoes fission with slow neutron capture; cobalt-60, because it is radioactive and has a convenient half-life and gamma-ray energy for industrial X-ray analysis; iodine-131, because it is radioactive and concentrates in the thyroid gland; and lead-208, because its nucleus has a special stability due to a favored arrangement of both protons and neutrons. Many other isotopes, including a few stable ones, are used as tags or tracers in chemical and biological experiments. Carbon-14 is used for dating organic matter.

History: In 1907, H. N. McCoy and W. H. Ross discovered that certain radioactive decay products have the same chemical properties as thorium but differ from thorium in atomic weight. This discovery provided the first evidence that some atoms have the same chemical properties but different atomic weights.

In 1913, Frederick Soddy introduced the term “isotopes” to designate substances that occupy the same place in the periodic table of the elements. That same year, Joseph J. Thomson concluded from his experiments that neon actually is a mixture of two gases, one with an atomic weight of about 20 and the other with an atomic weight of about 22. In 1919, Francis Aston used a mass spectrograph, which he had developed, to show conclusively that there are two isotopes of neon. In the next few years, Aston, Arthur Dempster, and others analyzed many elements and showed that most of them occur as mixtures of two or more isotopes.

The first artificial production of a radioisotope was achieved by Frédéric and Irène Joliot-Curie in 1934, when they bombarded aluminum with alpha particles and produced a radioactive isotope of phosphorus.

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