Atomic mass unit

The unified atomic mass unit or dalton (symbol: u, or Da or AMU) is a standard unit of mass that quantifies mass on an atomic or molecular scale (atomic mass). One unified atomic mass unit is approximately the mass of one nucleon (either a single proton or neutron) and is numerically equivalent to 1 g/mol.[1] It is defined as one twelfth of the mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state and at rest,[2] and has a value of 1.660539040(20)×10−27 kg,[3] or approximately 1.66 yoctograms.[4] The CIPM has categorised it as a non-SI unit accepted for use with the SI, and whose value in SI units must be obtained experimentally.[2]

The atomic mass unit (amu) without the "unified" prefix is technically an obsolete unit based on oxygen, which was replaced in 1961. However, many sources still use the term amu but now define it in the same way as u (i.e., based on carbon-12).[5][6] In this sense, most uses of the terms atomic mass units and amu, today, actually refer to unified atomic mass unit. For standardization, a specific atomic nucleus (carbon-12 vs. oxygen-16) had to be chosen because the average mass of a nucleon depends on the count of the nucleons in the atomic nucleus due to mass defect. This is also why the mass of a proton or neutron by itself is more than (and not equal to) 1 u.

The atomic mass unit is not the unit of mass in the atomic units system, which is rather the electron rest mass (me).

Until the 2019 redefinition of SI base units, the number of daltons in a gram is exactly the Avogadro number by definition, or equivalently, a dalton is exactly equivalent to 1 gram/mol. Thereafter, these relationships will no longer be exact, but they will still be extremely accurate approximations.[7][Note 1]

Unified atomic mass unit
Unit systemPhysical constant
(Accepted for use with the SI)
Unit ofmass
Symbolu or Da 
Named afterJohn Dalton
1 u or Da in ...... is equal to ...
   kg   1.660539040(20)×10−27
   MeV/c2   931.4940954(57)
   me   1822.888486192(53)

History of the atomic mass unit

The standard atomic weight (or atomic weight) scale has traditionally been a relative value, that is without a unit, with the first relative atomic mass basis suggested by John Dalton in 1803 as 1H.[8] Despite the initial mass of 1H being used as the natural unit for relative atomic mass, it was suggested by Wilhelm Ostwald that relative atomic mass would be best expressed in terms of units of 1/16 mass of oxygen (1903). This evaluation was made prior to the discovery of the existence of elemental isotopes, which occurred in 1912.[8]

The discovery of isotopic oxygen in 1929 led to a divergence in relative atomic mass representation, with isotopically weighted oxygen (i.e., naturally occurring oxygen relative atomic mass) given a value of exactly 16 atomic mass units (amu) in chemistry, while pure 16O (oxygen-16) was given the mass value of exactly 16 amu in physics.

The divergence of these values could result in errors in computations, and was unwieldy. The chemistry amu, based on the relative atomic mass (atomic weight) of natural oxygen (including the heavy naturally-occurring isotopes 17O and 18O), was about 1.000282 as massive as the physics amu, based on pure isotopic 16O.

For these and other reasons, the reference standard for both physics and chemistry was changed to carbon-12 in 1961.[9] The choice of carbon-12 was made to minimise further divergence with prior literature.[8] The new and current unit was referred to as the unified atomic mass unit, u.[10] and given a new symbol, "u", which replaced the now deprecated "amu" that had been connected to the old oxygen-based system. The dalton (Da) is another name for the unified atomic mass unit.[11]

Despite this change, modern sources often still use the old term "amu" but define it as u (1/12 of the mass of a carbon-12 atom), as mentioned in the article's introduction. Therefore, in general, "amu" likely does not refer to the old oxygen standard unit, unless the source material originates from the 1960s or before.


The unified atomic mass unit and the dalton are different names for the same unit of measure. As with other unit names such as watt and newton, dalton is not capitalized in English, but its symbol, Da, is capitalized. With the introduction of the name dalton, there has been a gradual change towards using that name in preference to the name, unified atomic mass unit:

  • In 1993, the International Union of Pure and Applied Chemistry (IUPAC) approved the use of the dalton with the qualification that the CGPM had not given its approval.[12]
  • In 2003, the Consultative Committee for Units, part of the CIPM, recommended a preference for the usage of the "dalton" over the "unified atomic mass unit" as it "is shorter and works better with prefixes".[13]
  • In 2005, the International Union of Pure and Applied Physics endorsed the use of the dalton as an alternative to the unified atomic mass unit.[14]
  • In 2006, in the 8th edition of the formal definition of SI, the CIPM cataloged the dalton alongside the unified atomic mass unit as a "Non-SI unit whose values in SI units must be obtained experimentally: Units accepted for use with the SI".[2] The definition also noted that "The dalton is often combined with SI prefixes ..."
  • In 2009, when the International Organization for Standardization published updated versions of ISO 80000, it gave mixed messages as to whether or not the unified atomic mass unit had been deprecated: ISO 80000-1:2009 (General), identified the dalton as having "earlier [been] called the unified atomic mass unit u",[15] but ISO 80000-10:2009 (atomic and nuclear physics) catalogued both as being alternatives for each other.[16]
  • The 2010 version of the Oxford University Press style guide for authors in life sciences gave the following guidance: "Use the Système international d'unités (SI) wherever possible ... The dalton (Da) or more conveniently the kDa is a permitted non-SI unit for molecular mass or mass of a particular band in a separating gel."[17] At the same time, the author guidelines for the journal "Rapid Communications in Mass Spectrometry" stated "The dalton (Da) is a unit of mass normally used for the molecular weight ... use of the Da in place of the u has become commonplace in the mass spectrometry literature ... The 'atomic mass unit', abbreviated 'amu', is an archaic unit".[18]
  • In 2012, in response to the proposed redefinition of the kilogram, it was proposed that the dalton be redefined as being 0.001/NA kg, thereby breaking the link with 12C. This would result in the dalton and the atomic mass unit having slightly different definitions, but the suggestion is that the older unit should be superseded by the "new" dalton.[19]
  • In 2018, the draft Ninth SI Brochure[7] associated with the 2019 redefinition of SI base units retains the definition of the Dalton and atomic mass unit unchanged in term of the mass of a carbon-12 atom.

Relationship to the International System of Units: SI

The former definition of the mole, an SI base unit, was accepted by the CGPM in 1971 as:

  1. The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12; its symbol is "mol".
  2. When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.

However, the first part of this definition will be changed on 20 May 2019 to:

One consequence of this change is that the current defined relationship between the mass of the 12C atom, the dalton, the kilogram, and the Avogadro number will no longer be valid. One of the following must change:

  • The mass of a 12C atom is exactly 12 dalton.
  • The number of dalton in a gram is exactly the numerical value of the Avogadro number: (i.e., 1 g/Da = 1 mol ⋅ NA).

The wording of the ninth SI Brochure[7][Note 1] implies that the first statement remains valid, which means that the second is no longer true. The molar mass constant, while still with great accuracy remaining 1 g/mol, is no longer exactly equal to that. Given that the unified atomic mass unit is one twelfth the mass of one atom of carbon-12, meaning the mass of such an atom is 12 u, it follows that there are only approximately NA atoms of carbon-12 in 0.012 kg of carbon-12. This can be expressed mathematically as


Molecular masses of proteins are often expressed in kilodaltons (kDa or kD). For example, a molecule of a protein with molar mass 64000 g⋅mol−1 has a mass of 64 kDa.[1]

In research and commerce, the degree of polymerization of synthetic polymers is conventionally expressed in daltons.

The US Supreme Court based a major precedent of appellate law on a disputed case of counting daltons for a molecular distribution.[22]


See also


  1. ^ a b A footnote in Table 8 on non-SI units states: "The dalton (Da) and the unified atomic mass unit (u) are alternative names (and symbols) for the same unit, equal to 1/12 of the mass of a free carbon 12 atom, at rest and in its ground state."


  1. ^ a b Berg, Jeremy M.; Tymoczko, John L.; Stryer, Lubert (2007). "2". Biochemistry (6th ed.). New York: Freeman. p. 35. ISBN 978-0-7167-8724-2.
  2. ^ a b c International Bureau of Weights and Measures (2006), The International System of Units (SI) (PDF) (8th ed.), p. 126, ISBN 92-822-2213-6, archived (PDF) from the original on 2017-08-14
  3. ^ "CODATA Value: atomic mass constant". The NIST Reference on Constants, Units, and Uncertainty. US National Institute of Standards and Technology. June 2015. Retrieved 2015-09-25. 2014 CODATA recommended values
  4. ^ Unified Atomic mass unit. Fundamental Physical Constants from NIST
  5. ^ Chang, Raymond (2005). Physical Chemistry for the Biosciences. p. 5. ISBN 978-1-891389-33-7.
  6. ^ Kelter, Paul B.; Mosher, Michael D.; Scott, Andrew (2008). Chemistry: The Practical Science. 10. p. 60. ISBN 978-0-547-05393-6.
  7. ^ a b c "Draft of the ninth SI Brochure" (PDF). BIPM. 5 February 2018. Retrieved 12 November 2018.
  8. ^ a b c Petley, B. W. (1989), "The atomic mass unit", IEEE Trans. Instrum. Meas., 38 (2): 175–79, doi:10.1109/19.192268
  9. ^ Holden, Norman E. (2004), "Atomic Weights and the International Committee—A Historical Review", Chem. Int., 26 (1): 4–7
  10. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "unified atomic mass unit". doi:10.1351/goldbook.U06554
  11. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "dalton". doi:10.1351/goldbook.D01514
  12. ^ Mills, Ian; Cvitaš, Tomislav; Homann, Klaus; Kallay, Nikola; Kuchitsu, Kozo (1993). Quantities, Units and Symbols in Physical Chemistry International Union of Pure and Applied Chemistry; Physical Chemistry Division (PDF) (2nd ed.). International Union of Pure and Applied Chemistry and published for them by Blackwell Science Ltd. ISBN 978-0-632-03583-0.
  13. ^ "Consultative Committee for Units (CCU); Report of the 15th meeting (17–18 April 2003) to the International Committee for Weights and Measures" (PDF). Retrieved 14 Aug 2010.
  14. ^ "IUPAP: C2: Report 2005". Retrieved 2018-07-15.
  15. ^ International Standard ISO 80000-1:2009 – Quantities and Units – Part 1: General, International Organization for Standardization, 2009
  16. ^ International Standard ISO 80000-10:2009 – Quantities and units – Part 10: Atomic and nuclear physics, International Organization for Standardization, 2009
  17. ^ "Instructions to Authors". AoB Plants. Oxford journals; Oxford University Press. Retrieved 2010-08-22.
  18. ^ "Author guidelines". Rapid Communications in Mass Spectrometry. Wiley-Blackwell. 2010.
  19. ^ Leonard, B P (2012). "Why the dalton should be redefined exactly in terms of the kilogram". Metrologia. 49 (4): 487–491. Bibcode:2012Metro..49..487L. doi:10.1088/0026-1394/49/4/487.
  20. ^ CIPM Report of 106th Meeting Archived 27 January 2018 at the Wayback Machine Retrieved 7 April 2018
  21. ^ "Redefining the Mole". NIST. NIST. 23 October 2018. Retrieved 24 October 2018.
  22. ^ "Supreme Court Opinion in 'Teva Pharmaceuticals USA, Inc. v. Sandoz, Inc.'" (PDF).
  23. ^ Opitz CA, Kulke M, Leake MC, Neagoe C, Hinssen H, Hajjar RJ, Linke WA (October 2003). "Damped elastic recoil of the titin spring in myofibrils of human myocardium". Proc. Natl. Acad. Sci. U.S.A. 100 (22): 12688–93. Bibcode:2003PNAS..10012688O. doi:10.1073/pnas.2133733100. PMC 240679. PMID 14563922.

External links


As an abbreviation, AMU may refer to:

Academy of Performing Arts in Prague, Prague, Czech republic

Adam Mickiewicz University, Poznań, Poland

Agency for the Modernisation of Ukraine

African Mathematical Union

Aix-Marseille University, Aix-en-Provence/Marseille, France

Aligarh Muslim University, Aligarh, India

American Military University, Charles Town, West Virginia

Arab Maghreb Union

Atomic mass unit, used to express atomic and molecular masses.

Ave Maria University, Florida, USA

Auxiliary Memory Units

African and Malagasy Union

Asian Monetary Unit

United States Army Medical Unit

Northern Rhodesian African Mineworkers' Union, a former Northern Rhodesian trade unionAmu may refer to:

Amu (film), Indian film starring Konkona Sen Sharma

Amu Hinamori, a fictional character from the manga series Shugo Chara! by Peach-Pit

Amu Darya, a river in central Asia

Amu (pharaoh)

Atomic mass

The atomic mass (ma) is the mass of an atom. Its unit is the unified atomic mass units (abbr. u) where 1 unified atomic mass unit is defined as ​1⁄12 of the mass of a single carbon-12 atom, at rest. For atoms, the protons and neutrons of the nucleus account for nearly all of the total mass, and the atomic mass measured in u has nearly the same value as the mass number.

When divided by unified atomic mass units, or daltons (abbr. Da), to form a pure numeric ratio, the atomic mass of an atom becomes a dimensionless value called the relative isotopic mass (see section below). Thus, the atomic mass of a carbon-12 atom is 12 u (or 12 Da), but the relative isotopic mass of a carbon-12 atom is simply 12.

The atomic mass or relative isotopic mass refers to the mass of a single particle, and therefore is tied to a certain specific isotope of an element. The dimensionless standard atomic weight instead refers to the average (mathematical mean) of atomic mass values of a typical naturally-occurring mixture of isotopes for a sample of an element. Atomic mass values are thus commonly reported to many more significant figures than atomic weights. Standard atomic weight is related to atomic mass by the abundance ranking of isotopes for each element. It is usually about the same value as the atomic mass of the most abundant isotope, other than what looks like (but is not actually) a rounding difference.

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 binding energy mass loss (as per E = mc2).

Atomic mass constant

In physics and chemistry, the atomic mass constant, mu, is one twelfth of the mass of an unbound atom of carbon-12 at rest and in its ground state. It serves to define the atomic mass unit and is, by definition, equal to 1 u. It is inverse of Avogadro constant (1/NA) when expressed in grams (instead of SI unit kilogram). The CODATA recommended value is 1.660539040(20)×10−27 kg.

In practice, the atomic mass constant is determined from the electron rest mass me and the electron relative atomic mass Ar(e) (that is, the mass of the electron on a scale where 12C = 12). The relative atomic mass of the electron can be measured in cyclotron experiments, while the rest mass of the electron can be derived from other physical constants.

where c is the speed of light, h is the Planck constant, α is the fine-structure constant, and R is the Rydberg constant.

The current (CODATA 2014) uncertainty in the value of the atomic mass constant – relative uncertainty 1.2×10−8 – is almost entirely due to the uncertainty in the value of the Planck constant in SI units. With the 2019 redefinition of SI base units, the relative uncertainty will improve to 4.7×10−10, which will be almost entirely due to the uncertainty in the fine-structure constant.

Dda (DNA-dependent ATPase)

Dda (short for DNA-dependent ATPase; also known as Dda helicase and Dda DNA helicase) is the 439-amino acid 49,897-atomic mass unit protein coded by the Dda gene of the bacteriophage T4 phage, a virus that infects enterobacteria.


ISOLTRAP is a tandem Penning trap mass spectrometer at the On-Line Isotope Mass Separator at CERN. The facility plays a leading role in the field of high precision mass spectrometry of radioactive ions. The masses of more than 200 short-lived nuclides have been measured with a relative uncertainty of typically dm/m ~ 1x10−7 and even almost up to one order of magnitude lower in some special cases.

Recently, the performance of the Penning trap mass spectrometer ISOLTRAP has been considerably enhanced. Major technical improvements were implemented to increase the range of accessible nuclei to those that are produced in minute quantities of only 100 ions/s and to nuclei with half-lives down to ~50 ms as well as to decrease the typical relative uncertainty down to ~1x10−8. In particular, a linear radiofrequency quadrupole (RFQ) trap and more recently a multi-reflection time-of-flight (MRTOF) mass spectrometer were added to the ISOLTRAP spectrometer. Since the unified atomic mass unit is defined as 1/12 of the mass of 12C the calibration of the magnetic field with carbon clusters allows absolute mass measurements. For this purpose, a laser-ablation source of carbon-cluster ions is installed at ISOLTRAP.

Accurate mass measurements of short-lived nuclides are of high interest for a number of reasons. From the measured atomic masses it is possible to compute nuclear binding energies, which are sensitive to nuclear structure effects like the location of shell and sub-shell closures, pairing, or the onset of deformation. In combination with a precise study of super-allowed beta emission they provide tests of the Standard Model. Additionally, masses of unstable nuclei are the most critical nuclear physics parameters for reliable nucleosynthesis calculations in astrophysics.

Isotopes of cobalt

Naturally occurring cobalt (27Co) is composed of 1 stable isotope, 59Co. 28 radioisotopes have been characterized with the most stable being 60Co with a half-life of 5.2714 years, 57Co with a half-life of 271.8 days, 56Co with a half-life of 77.27 days, and 58Co with a half-life of 70.86 days. All of the remaining radioactive isotopes have half-lives that are less than 18 hours and the majority of these have half-lives that are less than 1 second. This element also has 11 meta states, all of which have half-lives less than 15 minutes.

The isotopes of cobalt range in atomic weight from 47Co to 75Co. The primary decay mode for isotopes with atomic mass unit values less than that of the most abundant stable isotope, 59Co, is electron capture and the primary mode of decay for those of greater than 59 atomic mass units is beta decay. The primary decay products before 59Co are iron isotopes and the primary products after are nickel isotopes.

Radioactive isotopes can be produced by various nuclear reactions. For example, the isotope 57Co is produced by cyclotron irradiation of iron. The principal reaction involved is the (d,n) reaction 56Fe + 2H → n + 57Co.

Mass (mass spectrometry)

The mass recorded by a mass spectrometer can refer to different physical quantities depending on the characteristics of the instrument and the manner in which the mass spectrum is displayed.

Mass excess

The mass excess of a nuclide is the difference between its actual mass and its mass number in atomic mass units. It is one of the predominant methods for tabulating nuclear mass. The mass of an atomic nucleus is well approximated (less than 0.1% difference for most nuclides) by its mass number, which indicates that most of the mass of a nucleus arises from mass of its constituent protons and neutrons. Thus, the mass excess is an expression of the nuclear binding energy, relative to the binding energy per nucleon of carbon-12 (which defines the atomic mass unit). If the mass excess is negative, the nucleus has more binding energy than 12C, and vice versa. If a nucleus has a large excess of mass compared to a nearby nuclear species, it can radioactively decay, releasing energy.

Mass number

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 determines the atomic mass of atoms. Because protons and neutrons both are baryons, the mass number A is identical with the baryon number B as of the nucleus as of the whole atom or ion. The mass number is different for each different isotope of a chemical element. This is not the same as the atomic number (Z) which denotes the number of protons in a nucleus, and thus uniquely identifies an element. Hence, the difference between the mass number and the atomic number gives the number of neutrons (N) in a given nucleus: .

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 12
, 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: 12
. This is technically redundant, as each element is defined by its atomic number, so it is often omitted.

Milli mass unit

The milli mass unit or (mmu) is used as a unit of mass by some scientific authors even though this unit is not defined by the IUPAP red book nor by the IUPAC green book. It is a short form of the tongue-breaking but formally more correct "milli unified atomic mass unit" (mu) and equivalent to 1/1000 of the unified atomic mass unit (u). A more modern name is the millidalton (mDa)

since the "unified atomic mass unit" is more and more displaced by the unit dalton. (1 Da = 1 u)

Since 1961 the unified atomic mass unit "u" has been defined as 1/12 the mass of 12C. Before that the atomic mass unit "amu" was defined as 1/16 the mass of 16O (physics) and as 1/16 the mass of O (chemistry). Thus the publication date in literature ought to be heeded when reading about the milli mass unit as its name does not reveal whether it refers to the old amu or the newer u.

The mass excess is usually indicated in mu or mmu.

In mass spectrometry the mass accuracy of a mass analyzer is often indicated in mu, even though a more correct unit would be mTh since mass spectrometers measure the mass-to-charge ratio, not the mass. The relative mass accuracy is often indicated in ppm, even though this is no longer supported by the IUPAC green book which suggests using units like μTh/Th instead of ppm.

Mole (unit)

The mole is the base unit of amount of substance in the International System of Units (SI). Effective 20 May 2019, the mole is defined as the amount of a chemical substance that contains exactly 6.02214076×1023 (Avogadro constant) constitutive particles, e.g., atoms, molecules, ions or electrons.This definition was adopted in November 2018, revising its old definition based on the number of atoms in 12 grams of carbon-12 (12C) (the isotope of carbon with relative atomic mass 12 Da by definition). The mole is an SI base unit, with the unit symbol mol.

The mole is widely used in chemistry as a convenient way to express amounts of reactants and products of chemical reactions. For example, the chemical equation 2H2 + O2 → 2H2O can be interpreted to mean that 2 mol dihydrogen (H2) and 1 mol dioxygen (O2) react to form 2 mol water (H2O). The mole may also be used to represent the number of atoms, ions, or other entities in a given sample of a substance. The concentration of a solution is commonly expressed by its molarity, defined as the amount of dissolved substance per unit volume of solution, for which the unit typically used is moles per litre (mol/l), commonly abbreviated M.

The term gram-molecule was formerly used for essentially the same concept. The term gram-atom has been used for a related but distinct concept, namely a quantity of a substance that contains an Avogadro's number of atoms, whether isolated or combined in molecules. Thus, for example, 1 mole of MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2.

Molecular mass

Relative molecular mass or molecular weight is the mass of a molecule. It is calculated as the sum of the relative atomic masses of each constituent element multiplied by the number of atoms of that element in the molecular formula. The molecular mass of small to medium size molecules, measured by mass spectrometry, determines stoichiometry. For large molecules such as proteins, methods based on viscosity and light-scattering can be used to determine molecular mass when crystallographic data are not available.


A proton is a subatomic particle, symbol p or p+, with a positive electric charge of +1e elementary charge and a mass slightly less than that of a neutron. Protons and neutrons, each with masses of approximately one atomic mass unit, are collectively referred to as "nucleons".

One or more protons are present in the nucleus of every atom; they are a necessary part of the nucleus. The number of protons in the nucleus is the defining property of an element, and is referred to as the atomic number (represented by the symbol Z). Since each element has a unique number of protons, each element has its own unique atomic number.

The word proton is Greek for "first", and this name was given to the hydrogen nucleus by Ernest Rutherford in 1920. In previous years, Rutherford had discovered that the hydrogen nucleus (known to be the lightest nucleus) could be extracted from the nuclei of nitrogen by atomic collisions. Protons were therefore a candidate to be a fundamental particle, and hence a building block of nitrogen and all other heavier atomic nuclei.

In the modern Standard Model of particle physics, protons are hadrons, and like neutrons, the other nucleon (particles present in atomic nuclei), are composed of three quarks. Although protons were originally considered fundamental or elementary particles, they are now known to be composed of three valence quarks: two up quarks of charge +2/3e and one down quark of charge –1/3e. The rest masses of quarks contribute only about 1% of a proton's mass, however. The remainder of a proton's mass is due to quantum chromodynamics binding energy, which includes the kinetic energy of the quarks and the energy of the gluon fields that bind the quarks together. Because protons are not fundamental particles, they possess a physical size, though not a definite one; the root mean square charge radius of a proton is about 0.84–0.87 fm or 0.84×10−15 to 0.87×10−15 m.At sufficiently low temperatures, free protons will bind to electrons. However, the character of such bound protons does not change, and they remain protons. A fast proton moving through matter will slow by interactions with electrons and nuclei, until it is captured by the electron cloud of an atom. The result is a protonated atom, which is a chemical compound of hydrogen. In vacuum, when free electrons are present, a sufficiently slow proton may pick up a single free electron, becoming a neutral hydrogen atom, which is chemically a free radical. Such "free hydrogen atoms" tend to react chemically with many other types of atoms at sufficiently low energies. When free hydrogen atoms react with each other, they form neutral hydrogen molecules (H2), which are the most common molecular component of molecular clouds in interstellar space.

Relative atomic mass

Relative atomic mass (symbol: Ar) or atomic weight is a dimensionless physical quantity defined as the ratio of the average mass of atoms of a chemical element in a given sample to the atomic mass constant. The atomic mass constant (symbol: mu) is defined as being 1/12 of the mass of a carbon-12 atom. Since both quantities in the ratio are masses, the resulting value is dimensionless; hence the value is said to be relative.

For a single given sample, the relative atomic mass of a given element is the weighted arithmetic mean of the masses of the individual atoms (including their isotopes) that are present in the sample. This quantity can vary substantially between samples because the sample's origin (and therefore its radioactive history or diffusion history) may have produced unique combinations of isotopic abundances. For example, due to a different mixture of stable carbon-12 and carbon-13 isotopes, a sample of elemental carbon from volcanic methane will have a different relative atomic mass than one collected from plant or animal tissues.

The more common, and more specific quantity known as standard atomic weight (Ar, standard) is an application of the relative atomic mass values obtained from multiple different samples. It is sometimes interpreted as the expected range of the relative atomic mass values for the atoms of a given element from all terrestrial sources, with the various sources being taken from Earth. "Atomic weight" is often loosely and incorrectly used as a synonym for standard atomic weight (incorrectly because standard atomic weights are not from a single sample). Standard atomic weight is nevertheless the most widely published variant of relative atomic mass.

Additionally, the continued use of the term "atomic weight" (for any element) as opposed to "relative atomic mass" has attracted considerable controversy since at least the 1960s, mainly due to the technical difference between weight and mass in physics. Still, both terms are officially sanctioned by the IUPAC. The term "relative atomic mass" now seems to be replacing "atomic weight" as the preferred term, although the term "standard atomic weight" (as opposed to the more correct "standard relative atomic mass") continues to be used.


U (named u , plural ues) is the 21st letter and the fifth vowel in the ISO basic Latin alphabet. It is preceded by T, and is followed by V.

Base units
Derived units
with special names
Other accepted units
See also

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