Isotope | Research & Encyclopedia Articles

Isotope

The discovery of radioactivity by Henri Becquerel in 1896 opened a new field of research for chemists. An immediate and obvious question had to do with the materials that produced the radioactivity that Becquerel observed. In July, 1898, Marie Curie announced the discovery of polonium, which, along with uranium, produced some of the radiation observed by Becquerel. Five months later, she identified radium, a second new element that was also radioactive. Over the next two decades, claims for nearly 50 new radioactive elements were put forward.

These claims presented problems for chemists. Only a handful of empty spaces remained in the periodic table for new elements. They did not know how these new radioactive elements would be placed in the table?

Some possible answers to this dilemma appeared quite early. Chemists noted that many of the new "elements" seemed to be chemically indistinguishable from each other or from other known elements. For example, three supposedly different radioactive and gaseous elements were so much alike that they could not be distinguished from each other. In 1900, K. A. Hofmann (1870-1940) and E. Strauss showed that radium-D, one of the supposed new "elements," was chemically identical to lead. In fact, they were unable to separate the two from each other by any known chemical means. Also, the element discovered by Bertram Boltwood in 1907 and named by him ionium turned out to be chemically inseparable from thorium.

Another clue to the puzzle came from research that was being done onatomic weights. One postulate of John Dalton 's atomic theory had been that all atoms of an element were exactly identical in every respect. This postulate applied to atomic weight as well as any other property. Thus, Dalton assumed that all atoms of lead (as an example) had the same weight.

But research in the early 1900's showed that this postulate may not have been correct. For example, Theodore William Richards and M. E. Lembert found that the atomic weight of lead taken from radioactive ores ranged anywhere from 206.4 to 208.4.

A possible explanation of these facts was offered by the English chemist Frederick Soddy and, independently, by the Polish chemist Kasimir Fajans (1887-1975). Soddy observed that an atom that loses an alpha particle or a beta particle changes in a very predictable way into another kind of atom. An alpha particle has a mass of 4 and a charge of +2. When an atom loses an alpha particle, therefore, it must change into another atom with a positive charge two less than its original value. In terms of the periodic table, that means the new atom must occupy a space two places to the left of the original atom. An atom that loses a beta particle experiences no change in mass, but it gains a single positive charge. This means that the atom will occupy a space one place to the right of its original position. Soddy called these observations his radioactive displacement law.

One logical conclusion of the radioactive displacement law is that a single position in the periodic table might be occupied by two or more kinds of atoms. For example, an atom in position 90 could be produced by the alpha decay of element 92 or by the beta decay of element 89. These two pathways would result in an atom with the same atomic number (90), but two different atomic weights.

Soddy chose the name isotopes, meaning "the same place, " for these forms of an element. In modern terms, isotopes are forms of an element that have the same number of protons (and, hence, the same atomic number) but different numbers of neutrons (and, hence, different atomic weights).

Soddy's explanation clarified the situation with the 50 or so new "elements." It demonstrated that only five of the reported elements were actually such. The remaining substances were isotopes of these or of existing elements.

Although the concept of isotopes grew out of the study of radioactivity, its application to stable elements soon became apparent. In 1912, Joseph J. Thomson had observed that ordinary, stable neon appeared to exist in two forms. At first, he was unable to explain this observation. He thought that he had found a new compound of neon or even a new element mixed with the neon. He assigned the task of resolving this issue to a student, Francis Aston.

Aston was able to solve this problem with the aid of a new machine, the mass spectrograph, that he designed and built. He was able to show that the two forms of neon observed by Thomson were, in fact, isotopes of the element. The two isotopes had masses of 20 and 22. This evidence confirmed the fact that isotopes existed among stable elements and were not peculiar to radioactive elements.

The possible uses of isotopes as research tools became obvious soon after their discovery. In 1918, the Hungarian chemist György Hevesy used a radioactive isotope of lead to study the growth of plants. When lead was added to the water given to plants, Hevesy followed the radiation emitted by the lead, allowing him to trace its path through the plant's structures. Hevesy's work was the first in which radioactive isotopes were used as tracers.

In later years, scientists began to use isotopes of elements that occur naturally in plants and animals (lead does not), such as radioactive carbon, oxygen, nitrogen, and phosphorus. In 1935, the German biochemist Rudolf Schoenheimer (1898-1941) demonstrated that non-radioactive isotopes can also be used as tracers. He synthesized organic compounds containing deuterium and studied their movement through living systems. Since deuterium (hydrogen-2) is twice as heavy as normal hydrogen, it can easily be detected in an organism. The use of both radioactive and stable isotopes as tracers is now common in industry, medicine, and research.