Before the discovery of radioactivity, all atoms of a particular element were assumed to be identical.
Pure substances always had the same characteristic chemical and physical properties, no matter where the sample was taken from or how it was obtained. If the samples are indistinguishable, scientists reasoned, the atoms or molecules making up the samples should also be indistinguishable.
When evidence began to mount that an element's radioactive properties might vary from sample to sample, scientists were astonished. Thorium, for example, seemed to have at least two varieties. Thorium in naturally occurring minerals like thorite (ThSiO4) spits out alpha particles (helium nuclei). But thorium isolated from decaying uranium emits beta particles (electrons). Otherwise, the properties of the two forms of thorium are identical. This was an unsettling discovery, because it was in direct conflict with one of the foundation postulates of Dalton's atomic theory.
If the two forms have identical chemical properties, their electronic structures must also be identical. If both forms have the same number of electrons, then both must have the same number of protons in their nuclei. But something had to be different about them, or they wouldn't emit different types of radiation.
Experiments with mass spectrometry revealed that each of these element forms had its own distinctive mass. The element forms were called isotopes, from the Greek words for "same place", because all of these forms fit into the same place on the periodic table.
After neutrons were discovered, it was realized that
isotopes are substances that have same number of protons in their nuclei, but
different numbers of neutrons. Isotopes have the same atomic number but different masses and mass numbers.
|Isotopes have the same|
atomic number but
different mass numbers.
Isotopy isn't limited to radioactive elements.
Lighter elements almost always occur naturally as mixtures of isotopes. For example, hydrogen can occur in three isotopes:
||Hydrogen-1 (sometimes called protium) has one proton and no neutrons in its nucleus. (Notice that isotopes are named by separating the element name and the mass number with a dash).
||Hydrogen-2 (also called
deuterium) has one proton and one neutron in its nucleus. In a natural sample of hydrogen one in every 6000 H atoms is hydrogen-2. |
Hydrogen-3 (also called tritium) is a radioactive isotope with one proton and two neutrons per nucleus. Tritium is present in extremely low concentrations in natural hydrogen, but can be produced in large quantities by nuclear weapons.
The atomic weight given for an element is actually an average of the isotope masses, not the weight of an individual atom of the element. For example, 98.89% of all carbon atoms are carbon-12 (with 6 protons and 6 neutrons, and a mass of 12.00000). Almost all of the remainder are carbon-13 (with 6 protons and 7 neutrons, and a mass of 13.003354). The atomic weight is 12.01, which is 0.9889×12.00000 + 0.0111×13.003354.
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Author: Fred Senese email@example.com