Definition: Any of two or more forms of a chemical element, having the same number of protons in the nucleus, or the same atomic number, but having different numbers of neutrons in the nucleus, or different atomic weights. There are 275 isotopes of the 81 stable elements, in addition to over 800 radioactive isotopes, and every element has known isotopic forms. Isotopes of a single element possess almost identical properties.
Origin of Word: Greek
The different possible versions of each element are called isotopes(/ahy-suh-tohp/). Atoms of the same element can have different numbers of neutrons. For example, the most common isotope of hydrogen has no neutrons at all; there's also a hydrogen isotope called deuterium, with one neutron, and another, tritium, with two neutrons. Another example that you can take is of Carbon, there are a lot of carbon (C) atoms in the Universe. The normal ones, that exist in abundance on earth, are carbon-12. Those atoms have 6 neutrons. There are a few atoms that don't have 6. Those odd ones may have 7 or even 8 neutrons. Carbon-14 actually has 8 neutrons. C-14 is considered an isotope of the element carbon.
You might have noticed that some elements, like Chlorine (35.5 u), have fractional atomic masses instead of a natural number. Since we count relative atomic masses on the basis that one proton/one neutron equals one atomic weight unit and the number of protons and neutrons can only be a natural number. But still why do some elements have fractional atomic masses? The answer is isotopes. Atomic masses are calculated by figuring out the amounts of each type of atom and isotope there are in the Universe. For carbon, there are a lot of C-12, a couple of C-13, and a few C-14 atoms. When you average out all of the masses, you get a number that is a little bit higher than 12 (the weight of a C-12 atom). The average atomic mass for the element is actually 12.011. But this is simply rounded off to 12 in most cases for convenience in calculations.