Carbon-12 in the context of "Stable isotope"

Carbon (from Latin carbo 'coal') is a chemical element; it has symbol C and atomic number 6. It is nonmetallic and tetravalent—meaning that its atoms are able to form up to four covalent bonds due to its valence shell exhibiting 4 electrons. It belongs to group 14 of the periodic table. Carbon makes up about 0.025 percent of Earth's crust. Three isotopes occur naturally, C and C being stable, while C is a radionuclide, decaying with a half-life of 5,700 years. Carbon is one of the few elements known since antiquity.

Carbon is the 15th most abundant element in the Earth's crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. Carbon's abundance, its unique diversity of organic compounds, and its unusual ability to form polymers at the temperatures commonly encountered on Earth, enables this element to serve as a common element of all known life. It is the second most abundant element in the human body by mass (about 18.5%) after oxygen.

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:

Carbon (₆C) has 14 known isotopes, from
C to
C as well as
C, of which only
C and
C are stable. The longest-lived radioisotope is
C, with a half-life of 5700 years. This is also the only carbon radioisotope found in nature, as trace quantities are formed cosmogenically by the reaction
N + n →
C +
H. The most stable artificial radioisotope is
C, which has a half-life of 20.34 min. All other radioisotopes have half-lives under 20 seconds, most less than 200 milliseconds. Lighter isotopes exhibit beta-plus decay into isotopes of boron and heavier ones beta-minus decay into isotopes of nitrogen, though at the limits particle emission occurs as well.

The Eocene (IPA: /ˈiːəsiːn, ˈiːoʊ-/ EE-ə-seen, EE-oh-) is a geological epoch that lasted from about 56 to 33.9 million years ago (Ma). It is the second epoch of the Paleogene Period in the modern Cenozoic Era. The name "Eocene" comes from Ancient Greek ἠώς (ēṓs), meaning "dawn", and καινός (kainós), meaning "new", and refers to the "dawn" of modern ('new') fauna that appeared during the epoch.

The Eocene spans the time from the end of the Paleocene Epoch to the beginning of the Oligocene Epoch. The start of the Eocene is marked by a brief period in which the concentration of the carbon isotope C in the atmosphere was exceptionally low in comparison with the more common isotope C. The average temperature of Earth at the beginning of the Eocene was about 27 degrees Celsius. The end is set at a major extinction event called the Grande Coupure (the "Great Break" in continuity) or the Eocene–Oligocene extinction event, which may be related to the impact of one or more large bolides in Siberia and in what is now Chesapeake Bay. As with other geologic periods, the strata that define the start and end of the epoch are well identified, though their exact dates are slightly uncertain.

Carbon-14, C-14, C or radiocarbon, is a radioactive isotope of carbon with an atomic nucleus containing 6 protons and 8 neutrons. Its presence in organic matter is the basis of the radiocarbon dating method pioneered by Willard Libby and colleagues (1949) to date archaeological, geological and hydrogeological samples. Carbon-14 was discovered on February 27, 1940, by Martin Kamen and Sam Ruben at the University of California Radiation Laboratory in Berkeley, California. Its existence had been suggested by Franz Kurie in 1934.

There are three naturally occurring isotopes of carbon on Earth: carbon-12 (C), which makes up 99% of all carbon on Earth; carbon-13 (C), which makes up 1%; and carbon-14 (C), which occurs in trace amounts, making up about 1.2 atoms per 10 atoms of carbon in the atmosphere. C and C are both stable; C is unstable, with half-life 5700±30 years, decaying into nitrogen-14 (
N) through beta decay. Pure carbon-14 would have a specific activity of 62.4 mCi/mmol (2.31 GBq/mmol), or 164.9 GBq/g. The primary natural source of carbon-14 on Earth is cosmic ray action on nitrogen in the atmosphere, and it is therefore a cosmogenic nuclide. Open-air nuclear testing between 1955 and 1980 contributed to this pool, however.

The dalton or unified atomic mass unit (symbols: Da or u, respectively) is a unit of mass defined as ⁠1/12⁠ of the mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state and at rest. It is a non-SI unit accepted for use with SI. The word "unified" emphasizes that the definition was accepted by both IUPAP and IUPAC. The atomic mass constant, denoted m_u, is an atomic-scale reference mass, defined identically, but it is not a unit of mass. Expressed in terms of m_a(C), the atomic mass of carbon-12: m_u = m_a(C)/12 = 1 Da. The dalton's numerical value in terms of the fixed-h kilogram is an experimentally determined quantity that, along with its inherent uncertainty, is updated periodically. As listed in the 9th edition, version 3.02, of the SI Brochure, the 2022 CODATA recommended value of the atomic mass constant expressed in the SI base unit kilogram is:

The previous value given for the dalton (1 Da = 1 u = m_u) was the 2018 CODATA recommended value:

The mole (symbol mol) is a unit of measurement, the base unit in the International System of Units (SI) for amount of substance, an SI base quantity proportional to the number of elementary entities of a substance. One mole is an aggregate of exactly 6.02214076×10 elementary entities (approximately 602 sextillion or 602 billion times a trillion), which can be atoms, molecules, ions, ion pairs, or other particles. The number of particles in a mole is the Avogadro number (symbol N₀) and the numerical value of the Avogadro constant (symbol N_A) has units of mol. The relationship between the mole, Avogadro number, and Avogadro constant can be expressed in the following equation:$${\displaystyle 1{\text{ mol}}={\frac {N_{0}}{N_{\text{A}}}}={\frac {6.02214076\times 10^{23}}{N_{\text{A}}}}}$$The current SI value of the mole is based on the historical definition of the mole as the amount of substance that corresponds to the number of atoms in 12 grams of C, which made the molar mass of a compound in grams per mole, numerically equal to the average molecular mass or formula mass of the compound expressed in daltons. With the 2019 revision of the SI, the numerical equivalence is now only approximate, but may still be assumed with high accuracy.

Conceptually, the mole is similar to the concept of dozen or other convenient grouping used to discuss collections of identical objects. Because laboratory-scale objects contain a vast number of tiny atoms, the number of entities in the grouping must be huge to be useful for work.

The mass number (symbol A, from the German word: Atomgewicht, "atomic weight"), also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. It is approximately equal to the atomic (also known as isotopic) mass of the atom expressed in daltons. Since protons and neutrons are both baryons, the mass number A is identical with the baryon number B of the nucleus (and also of the whole atom or ion). The mass number is different for each isotope of a given chemical element, and the difference between the mass number and the atomic number Z gives the number of neutrons (N) in the nucleus: N = A − Z.

The mass number is written either after the element name or as a superscript to the left of an element's symbol. For example, the most common isotope of carbon is carbon-12, or
C, which has 6 protons and 6 neutrons. The full isotope symbol would also have the atomic number (Z) as a subscript to the left of the element symbol directly below the mass number:
₆C.