Isotope

Isotopes are distinct nuclear species (or nuclides) of the same chemical element. They have the same atomic number (number of protons in their nuclei) and position in the periodic table (and hence belong to the same chemical element), but different nucleon numbers (mass numbers) due to different numbers of neutrons in their nuclei. While all isotopes of a given element have virtually the same chemical properties, they have different atomic masses and physical properties.

The term isotope comes from the Greek roots isos (ἴσος "equal") and topos (τόπος "place"), meaning "the same place": different isotopes of an element occupy the same place on the periodic table. It was coined by Scottish doctor and writer Margaret Todd in a 1913 suggestion to the British chemist Frederick Soddy, who popularized the term.

The number of protons within the atom's nucleus is called its atomic number and is equal to the number of electrons in the neutral (non-ionized) atom. Each atomic number identifies a specific element, but not the isotope; an atom of a given element may have a wide range in its number of neutrons. The number of nucleons (both protons and neutrons) in the nucleus is the atom's mass number, and each isotope of a given element has a different mass number.

For example, carbon-12, carbon-13, and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, and 14, respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons so that the neutron numbers of these isotopes are 6, 7, and 8 respectively.

A nuclide is a species of an atom with a specific number of protons and neutrons in the nucleus, for example, carbon-13 with 6 protons and 7 neutrons. Thus the terms are roughly synonymous, but the nuclide concept (referring to individual nuclear species) emphasizes nuclear properties over chemical properties, whereas the isotope concept (grouping all atoms of each element) emphasizes chemical over nuclear. The neutron number greatly affects nuclear properties, but its effect on chemical properties is negligible for most elements. Even for the lightest elements, whose ratio of neutron number to atomic number varies the most between isotopes, it usually has only a small effect although it matters in some circumstances (for hydrogen, the lightest element, the isotope effect is large enough to affect biology strongly). The term isotopes (originally also isotopic elements, now sometimes isotopic nuclides) is intended to imply comparison (like synonyms or isomers). For example, the nuclides ¹²
₆C, ¹³
₆C, ¹⁴
₆C are isotopes (nuclides with the same atomic number but different mass numbers), but ⁴⁰
₁₈Ar, ⁴⁰
₁₉K, ⁴⁰
₂₀Ca are isobars (nuclides with the same mass number). As the older and better-known term, isotope is however still used in some contexts where nuclide might be more appropriate, such as in nuclear technology and nuclear medicine.

An isotope/nuclide is specified by the name of the element (this indicates the atomic number) followed by a hyphen and the mass number (e.g. helium-3, helium-4, carbon-12, carbon-14, uranium-235 and uranium-239). When a chemical symbol is used, e.g. "C" for carbon, standard notation (also known as "AZE notation" as it is written ^A_ZE where A is the mass number, Z the atomic number, and E the element name) is to indicate the mass number (number of nucleons) with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e.g. ³
₂He, ⁴
₂He, ¹²
₆C, ¹⁴
₆C, ²³⁵
₉₂U, and ²³⁹
₉₂U). Because the atomic number is already fixed by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript (e.g. ³He, ⁴He, ¹²C, ¹⁴C, ²³⁵U, and ²³⁹U). The letter m (for metastable) is appended after the mass number to indicate a nuclear isomer, a metastable or energetically excited nuclear state (as opposed to the lowest-energy ground state), for example ^180m
₇₃Ta (tantalum-180m); a number can be appended to it to distinguish different metastable states, though this is rare in practice.

The common pronunciation of the AZE notation is different from how it is written: ⁴
₂He is commonly pronounced helium-four instead of four-two-helium, and ²³⁵
₉₂U uranium two-thirty-five (American English) or uranium-two-three-five (British) instead of 235-92-uranium or 235-uranium. This is not an error but the original spoken usage for isotope names, originating before AZE notation became established.

Some isotopes/nuclides are radioactive, and are therefore called radioisotopes or radionuclides, whereas others have never been observed to decay radioactively and are called stable isotopes or stable nuclides. For example, ¹⁴C is a radioactive form of carbon, while ¹²C and ¹³C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, of which 286 are primordial nuclides, meaning that they have existed since the Solar System's formation.