kb/data/en.wikipedia.org/wiki/Atomic_mass-0.md

title	chunk	source	category	tags	date_saved	instance
Atomic mass	1/3	https://en.wikipedia.org/wiki/Atomic_mass	reference	science, encyclopedia	2026-05-05T10:51:53.116357+00:00	kb-cron

Atomic mass (ma or m) is the mass of a single atom. The atomic mass mostly comes from the combined mass of the protons and neutrons in the nucleus, with minor contributions from the electrons and nuclear binding energy. The atomic mass of atoms, ions, or atomic nuclei is slightly less than the sum of the masses of their constituent protons, neutrons, and electrons, due to mass defect (explained by mass–energy equivalence: E = mc2). Atomic mass is often measured in dalton (Da) (a.k.a. unified atomic mass unit (u)). One dalton is equal to ⁠+1/12⁠ the mass of a carbon-12 atom in its natural state, given by the atomic mass constant mu = m(12C)/12 = 1 Da, where m(12C) is the atomic mass of carbon-12. Thus, the numerical value of the atomic mass of a nuclide when expressed in daltons is close to its mass number. The relative isotopic mass (see section below) can be obtained by dividing the atomic mass ma of an isotope by the atomic mass constant mu, yielding a dimensionless value. Thus, the atomic mass of a carbon-12 atom m(12C) is 12 Da by definition, but the relative isotopic mass of a carbon-12 atom Ar(12C) is simply 12. The sum of relative isotopic masses of all atoms in a molecule is the relative molecular mass. The atomic mass of an isotope and the relative isotopic mass refers to a certain specific isotope of an element. Because substances are usually not isotopically pure, it is convenient to use the elemental atomic mass which is the average atomic mass of an element, weighted by the abundance of the isotopes. The dimensionless (standard) atomic weight is the weighted mean relative isotopic mass of a (typical naturally occurring) mixture of isotopes.

== Relative isotopic mass == Relative isotopic mass (a property of a single atom) is not to be confused with the averaged quantity atomic weight (see above), that is an average of values for many atoms in a given sample of a chemical element. While atomic mass is an absolute mass, relative isotopic mass is a dimensionless number with no units. This loss of units results from the use of a scaling ratio with respect to a carbon-12 standard, and the word "relative" in the term "relative isotopic mass" refers to this scaling relative to carbon-12. The relative isotopic mass, then, is the mass of a given isotope (specifically, any single nuclide), when this value is scaled by the mass of carbon-12, where the latter has to be determined experimentally. Equivalently, the relative isotopic mass of an isotope or nuclide is the mass of the isotope relative to ⁠+1/12⁠ of the mass of a carbon-12 atom. For example, the relative isotopic mass of a carbon-12 atom is exactly 12. For comparison, the atomic mass of a carbon-12 atom is exactly 12 daltons. Alternately, the atomic mass of a carbon-12 atom may be expressed in any other mass units: for example, the atomic mass of a carbon-12 atom is 1.99264688270(62)×10−26 kg. As is the case for the related atomic mass when expressed in daltons, the relative isotopic mass numbers of nuclides other than carbon-12 are not whole numbers, but are always close to whole numbers. This is discussed fully below.

== Similar terms for different quantities == The atomic mass or relative isotopic mass are sometimes confused, or incorrectly used, as synonyms of relative atomic mass (also known as atomic weight) or the standard atomic weight (a particular variety of atomic weight, in the sense that it is standardized). However, as noted in the introduction, atomic mass is an absolute mass while all other terms are dimensionless. Relative atomic mass and standard atomic weight represent terms for (abundance-weighted) averages of relative atomic masses in elemental samples, not for single nuclides. Relative atomic mass and standard atomic weight will only be the same as the relative isotopic mass for elements with one stable isotope. The atomic mass (relative isotopic mass) is defined as the mass of a single atom, which is only one isotope (nuclide), and is not an abundance-weighted average, as in the case of relative atomic mass/atomic weight. The atomic mass or relative isotopic mass of each isotope and nuclide of a chemical element is, therefore, a number that can be measured to high precision, since every specimen of such a nuclide is identical in mass to every other specimen, as all atoms of a given type in the same energy state, and every specimen of a particular nuclide, are identical in mass. For example, every atom of oxygen-16 has exactly the same atomic mass (relative isotopic mass). For elements that have one naturally occurring isotope (mononuclidic elements) or one dominant isotope, the difference between the atomic mass of the most common isotope, and the (standard) relative atomic mass or (standard) atomic weight will be nil, or small enough to not effect most calculations. For elements with more than one isotope in significant quantities, the mass of individual atoms will be different from an average across all the isotopes. For elements that have more than one common isotope, the numerical difference in relative atomic mass (atomic weight) from even the most common relative isotopic mass, can be half a mass unit or more (e.g. see the case of chlorine where atomic weight and standard atomic weight are about 35.45). The atomic mass (relative isotopic mass) of an uncommon isotope can differ from the relative atomic mass, atomic weight, or standard atomic weight, by several mass units. Relative isotopic masses are always close to whole-number values, but never (except in the case of carbon-12) exactly a whole number, for two reasons: