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2.9: Isotopes
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2.9: Isotopes

Elements have a set number of protons that determines their atomic number. For example, all atoms with eight protons are oxygen. However, the number of neutrons can vary for an atom of the same element. These variations of elements, with the same number of protons but different numbers of neutrons, are called isotopes.

The mass number is the sum of protons and neutrons. Hence, isotopes of an element have the same atomic number, but different mass numbers. An element’s atomic mass, or atomic weight, is a weighted average of the masses of the element’s isotopes. The average is said to be weighted because it reflects the relative abundance of the different isotopes in the sample. In other words, the masses of the most common isotopes contribute most strongly to the average.

Several elements exist as multiple isotopes in nature, including carbon, potassium, and uranium. On the periodic table, the atomic mass of an element reflects the relative abundance of their naturally-occurring isotopes on Earth.

Isotopes are often discussed in the context of radioactivity. A radioactive element is essentially an element with an unstable nucleus. Most radioactive elements have atomic numbers of 84 or higher. Other elements have isotopes that are non-radioactive and, in most cases, at least one radioactive isotope, a radioisotope.

To become more stable, radioisotopes release subatomic particles. In the process, known as radioactive decay, they emit energy known as radiation. Radioactive decay can alter the number of protons in an element, effectively changing its identity.

Radiation can be used to help determine the age or thickness of different materials. In medicine, it is applied to diagnose and track medical conditions using PET scanners, as well as to treat cancer.


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