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The atomic number corresponds to the number of protons in an atom. Thus, isotopes of a particular element contain the same number of protons. The difference in atomic weights results from differences in the number of neutrons in the atomic nuclei. In scientific nomenclature, isotopes are specified by the name of the particular element by a hyphen and the number of nucleons (protons and neutrons) in the atomic nucleus (e.g., helium-3, carbon-12, carbon-14, iron-57, uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g., ³He, ¹²C, ¹⁴C, ⁵⁷Fe, ²³⁸U).

Collectively, the isotopes of the elements form the set of nuclides. A nuclide is a particular type of nucleus (characterised by A and Z). Strictly speaking, we should say that an element such as fluorine consists of one nuclide rather than that it has one isotope. Similarly, the tables at the foot of this article are tables of nuclides.

In a neutral atom, the number of electrons equals the number of protons. Thus, isotopes of a given element also have the same number of electrons and the same electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The primary exception is that, due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This "mass effect" is most pronounced for protium (¹H) and deuteriumDeuterium (symbol 2H is a stable isotope of hydrogen with a natural abundance of one part in 7000 of hydrogen. The nucleus of deuterium (called a deuteron has one proton and one neutron, whereas a normal hydrogen nucleus just has one proton. Deuterium is (the common name of ²H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion and stabilizing the nucleus. For this reason neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus, for example, although the neutron/proton ratio of ³He is 1/2, the neutron/proton ratio of ²³⁸U is >3/2. However, if too many neutrons are present, the nucleus becomes unstable.

Because isotopes of a given element have different numbers of neutrons they also have different neutron/proton ratios. This affects the nuclear stability, with the result that some isotopes are subject to nuclear decay. The decay of these radioactive isotopes (radioisotopes for short) is an important topic in nuclear physicsNuclear physics is that branch of physics concerned with the nucleus of the atom. Topics include: Strong interaction Radioactivity Models of the nucleus liquid drop model shell model interacting boson model Fission Fusion Nuclear reactions Applications Nu. By studying the manner in which this decay occurs, physicists gain insight into the properties of the atomic nucleus.

In general, several isotopes of each element can be found in nature. Stable isotopes are by far the most abundant; however, significant quantities of long-lived unstable isotopes, such as uranium-238, can also be found. Small amounts of short-lived radioactive isotopes are also present in nature. These arise as products of the decay of larger long-lived radioactive nuclei. The atomic mass for an element in the periodic table is the average of the natural abundanceNatural abundance refers to the prevalence of different isotopes of an element as found in nature. The weighted (by natural abundance) average mass of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an el of the isotopes of that element.

The amounts of the various isotopes on earth is ultimately the result of the amounts formed in stars and supernovae, and the subsequent decay patterns of the radioactive nuclei formed in these processes. After that, the formation of the solar systemA generic solar system (or planetary system consists of at least one star and various orbiting objects (such as asteroids, comets, moons, and planets). The term originated to describe the planetary system around Sol, the Latin name for our sun. The planet also influenced heavily on the proportions of different isotopes found here, since lighter nuclei was more easily blown away towards the outer parts of the solar system, by the solar wind immediately after the sunThe Sun (also called Sol is the star in our solar system. Planet Earth orbits the Sun. Other bodies that orbit the Sun include other planets, asteroids, meteoroids, comets and dust. Not all objects passing through the solar system have been orbitally capt was formed. This is also why the gas giants are located further from the sun.