Stable isotope in the context of Krypton-86

In geochemistry, paleoclimatology, and paleoceanography δC (pronounced "delta thirteen c") is an isotopic signature, a measure of the ratio of the two stable isotopes of carbon—C and C—reported in parts per thousand (per mil, ‰). The measure is also widely used in archaeology for the reconstruction of past diets, particularly to see if marine foods or certain types of plants were consumed.

The definition is, in per mille:

Cosmogenic nuclides (or cosmogenic isotopes) are rare nuclides (isotopes) created when a high-energy cosmic ray interacts with the nucleus of an in situ Solar System atom, causing nucleons (protons and neutrons) to be expelled from the atom (see cosmic ray spallation). These nuclides are produced within Earth materials such as rocks or soil, in Earth's atmosphere, and in extraterrestrial items such as meteoroids. By measuring cosmogenic nuclides, scientists are able to gain insight into a range of geological and astronomical processes. There are both radioactive and stable cosmogenic nuclides. Some of these radionuclides are tritium, carbon-14 and phosphorus-32.

Certain light (low atomic number) primordial nuclides (isotopes of lithium, beryllium and boron) are thought to have been created not only during the Big Bang, but also (and perhaps primarily) to have been made after the Big Bang, but before the condensation of the Solar System, by the process of cosmic ray spallation on interstellar gas and dust. This explains their higher abundance in cosmic dust as compared with their abundances on Earth. This also explains the overabundance of the early transition metals just before iron in the periodic table – the cosmic-ray spallation of iron produces scandium through chromium on the one hand and helium through boron on the other. However, the arbitrary defining qualification for cosmogenic nuclides of being formed "in situ in the Solar System" (meaning inside an already aggregated piece of the Solar System) prevents primordial nuclides formed by cosmic ray spallation before the formation of the Solar System from being termed "cosmogenic nuclides"—even though the mechanism for their formation is exactly the same. These same nuclides still arrive on Earth in small amounts in cosmic rays, and are formed in meteoroids, in the atmosphere, on Earth, "cosmogenically". However, beryllium (all of it stable beryllium-9) is present primordially in the Solar System in much larger amounts, having existed prior to the condensation of the Solar System, and thus present in the materials from which the Solar System formed.

Iridium is a chemical element; it has symbol Ir and atomic number 77. This very hard, brittle, silvery-white transition metal of the platinum group, is considered the second-densest naturally occurring metal (after osmium) with a density of 22.56 g/cm (0.815 lb/cu in) as defined by experimental X-ray crystallography. Ir and Ir are the only two naturally occurring isotopes of iridium, as well as the only stable isotopes; the latter is the more abundant. It is one of the most corrosion-resistant metals, even at temperatures as high as 2,000 °C (3,630 °F).

Iridium was discovered in 1803 in the acid-insoluble residues of platinum ores by the British chemist Smithson Tennant. The name iridium, derived from the Greek word iris (rainbow), refers to the various colors of its compounds. Iridium is one of the rarest elements in Earth's crust, with an estimated annual production of only 6,800 kilograms (15,000 lb) in 2023.

Carbon-12 (C) is the most abundant of the two stable isotopes of carbon (carbon-13 being the other), amounting to 98.93% of element carbon on Earth; its abundance is due to the triple-alpha process by which it is created in stars. Carbon-12 is of particular importance in its use as the standard from which atomic masses of all nuclides are measured, thus, its atomic mass is exactly 12 daltons by definition. Carbon-12 is composed of 6 protons, 6 neutrons, and 6 electrons.

See carbon-13 for means of separating the two isotopes, thereby enriching both.

Carbon-13 (C) is a natural, stable isotope of carbon with a nucleus containing six protons and seven neutrons. It constitutes about 1.07% of natural carbon and is one of the so-called environmental isotopes.

Helium-4 (
He) is a stable isotope of the element helium. It is by far the more abundant of the two naturally occurring isotopes of helium, making up virtually all the helium on Earth. Its nucleus consists of two protons and two neutrons and is identical to an alpha particle.

Helium-3 (He see also helion) is a light, stable isotope of helium with two protons and one neutron. (In contrast, the most common isotope, helium-4, has two protons and two neutrons.) Helium-3 and hydrogen-1 are the only stable nuclides with more protons than neutrons. It was discovered in 1939. Helium-3 atoms are fermionic and become a superfluid at the temperature of 2.491 mK.

Helium-3 occurs as a primordial nuclide, escaping from Earth's crust into its atmosphere and into outer space over millions of years. It is also thought to be a natural nucleogenic and cosmogenic nuclide, one produced when lithium is bombarded by natural neutrons, which can be released by spontaneous fission and by nuclear reactions with cosmic rays. Some found in the terrestrial atmosphere is a remnant of atmospheric and underwater nuclear weapons testing.

Argon (₁₈Ar) has 26 known isotopes, from Ar to Ar, of which three are stable (Ar, Ar, and Ar). On Earth, Ar makes up 99.6% of natural argon. The longest-lived radioactive isotopes are Ar with a half-life of 302 years, Ar with a half-life of 32.9 years, and Ar with a half-life of 35.01 days. All other isotopes have half-lives of less than two hours, and most less than one minute. Isotopes lighter than Ar decay to chlorine or lighter elements, while heavier ones beta decay to potassium.

The naturally occurring K, with a half-life of 1.248×10 years, decays to stable Ar by electron capture (10.72%) and by positron emission (0.001%), and also to stable Ca via beta decay (89.28%). These properties and ratios are used to determine the age of rocks through potassium–argon dating.

Uranium (₉₂U) is a naturally occurring radioactive element (radioelement) with no stable isotopes. It has two primordial isotopes, uranium-238 and uranium-235, that have long half-lives and are found in appreciable quantity in Earth's crust. The decay product uranium-234 is also found. Other isotopes such as uranium-233 have been produced in breeder reactors. In addition to isotopes found in nature or nuclear reactors, many isotopes with far shorter half-lives have been produced, ranging from U to U (except for U). The standard atomic weight of natural uranium is 238.02891(3).

Natural uranium consists of three main isotopes, U (99.2739–99.2752% natural abundance), U (0.7198–0.7202%), and U (0.0050–0.0059%). All three isotopes are radioactive (i.e., they are radioisotopes), and the most abundant and stable is uranium-238, with a half-life of 4.463×10 years (about the age of the Earth).

The Paleocene–Eocene thermal maximum (PETM), alternatively "Eocene thermal maximum 1 (ETM1)" and formerly known as the "Initial Eocene" or "Late Paleocene thermal maximum", was a geologically brief time interval characterized by a 5–8 °C (9–14 °F) global average temperature rise and massive input of carbon into the ocean and atmosphere. The event began, now formally codified, at the precise time boundary between the Paleocene and Eocene geological epochs. The exact age and duration of the PETM remain uncertain, but it occurred around 55.8 million years ago (Ma) and lasted about 200 thousand years (Ka).

The PETM arguably represents our best past analogue for which to understand how global warming and the carbon cycle operate in a greenhouse world. The time interval is marked by a prominent negative excursion in carbon stable isotope (δC) records from around the globe; more specifically, a large decrease in the C/C ratio of marine and terrestrial carbonates and organic carbon has been found and correlated across hundreds of locations. The magnitude and timing of the PETM (δC) excursion, which attest to the massive past carbon release to our ocean and atmosphere, and the source of this carbon remain topics of considerable current geoscience research.

Isotope geochemistry is an aspect of geology based upon the study of natural variations in the relative abundances of isotopes of various elements. Variations in isotopic abundance are measured by isotope-ratio mass spectrometry, and can reveal information about the ages and origins of rock, air or water bodies, or processes of mixing between them.

Stable isotope geochemistry is largely concerned with isotopic variations arising from mass-dependent isotope fractionation, whereas radiogenic isotope geochemistry is concerned with the products of natural radioactivity.

A monoisotopic element is an element which has one and only one stable isotope (nuclide). There are 26 such elements, listed below.

Stability is experimentally defined for chemical elements, as all nuclides with atomic numbers over 40 or 66 (depending on definition, see stable nuclide) are theoretically unstable, but apparently have half-lives so long that their decay has not been observed either directly or indirectly (from measurement of products).

A helion (symbol h) is the nucleus of a helium atom, a doubly positively charged cation. The term helion is a portmanteau of helium and ion, and in practice refers specifically to the nucleus of the helium-3 isotope, consisting of two protons and one neutron. The nucleus of the other (and far more common) stable isotope of helium, helium-4, consisting of two protons and two neutrons, is called an alpha particle or an alpha for short.

This particle is the daughter product in the beta-minus decay of tritium, an isotope of hydrogen:

Oganesson (₁₁₈Og) is a synthetic element created in particle accelerators, and thus a standard atomic weight cannot be given. Like all synthetic elements, it has no stable isotopes. The first and only isotope to be synthesized was Og in 2002 and 2005; it has a half-life of 0.7 milliseconds.

Bismuth (₈₃Bi) has 41 known isotopes, ranging from Bi to Bi. Bismuth has no stable isotopes, but does have one naturally occurring, very long-lived isotope; thus, the standard atomic weight can be given from that isotope, bismuth-209. Though it is now known to be radioactive, it may still be considered practically stable because it has a half-life of 2.01×10 years, which is more than a billion times the age of the universe.

Besides Bi, the most stable bismuth radioisotopes are Bi with a half-life of 3.04 million years, Bi with a half-life of 368,000 years and Bi, with a half-life of 31.22 years, none of which occur in nature. All other isotopes have half-lives under 15 days, most under two hours. Of naturally occurring radioisotopes, the most stable is radiogenic Bi with a half-life of 5.012 days. Bi is unusual for being a nuclear isomer with a half-life many orders of magnitude longer than that of the ground state.

The nuclear drip line is the boundary beyond which atomic nuclei are unbound with respect to the emission of a proton or neutron.

An arbitrary combination of protons and neutrons does not necessarily yield a stable nucleus. One can think of moving up or to the right across the table of nuclides by adding a proton or a neutron, respectively, to a given nucleus. However, adding nucleons one at a time to a given nucleus will eventually lead to a newly formed nucleus that immediately decays by emitting a proton (or neutron). Colloquially speaking, the nucleon has leaked or dripped out of the nucleus, hence giving rise to the term drip line.