The dalton or unified atomic mass unit (symbols: Da or u) is a unit of mass widely used in physics and chemistry. It is 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. The atomic mass constant, denoted ''m''u is defined identically, giving . This unit is commonly used in physics and chemistry to express the mass of atomic-scale objects, such as atoms, molecules, and elementary particles, both for discrete instances and multiple types of ensemble averages. For example, an atom of helium-4 has a mass of . This is an intrinsic property of the isotope and all helium-4 have the same mass. Acetylsalicylic acid (aspirin), , has an average mass of approximately . However, there are no acetylsalicylic acid molecules with this mass. The two most common masses of individual acetylsalicylic acid molecules are and . The molecular masses of proteins, nucleic acids, and other large polymers are often expressed with the units kilodaltons (kDa), megadaltons (MDa), etc. Titin, one of the largest known proteins, has a molecular mass of between 3 and 3.7 megadaltons. The DNA of chromosome 1 in the human genome has about 249 million base pairs, each with an average mass of about , or total. The mole is a unit of amount of substance, widely used in chemistry and physics, which was originally defined so that the mass of one mole of a substance, measured in grams, would be numerically equal to the average mass of one of its constituent particles, measured in daltons. That is, the molar mass of a chemical compound was meant to be numerically equal to its average molecular mass. For example, the average mass of one molecule of water is about 18.0153 daltons, and one mole of water is about 18.0153 grams. A protein whose molecule has an average mass of would have a molar mass of . However, while this equality can be assumed for almost all practical purposes, it is now only approximate, because of the way mole was redefined on 20 May 2019. In general, the mass in daltons of an atom is numerically close, but not exactly equal to the number of nucleons ''A'' contained in its nucleus. It follows that the molar mass of a compound (grams per mole) is numerically close to the average number of nucleons contained in each molecule. By definition, the mass of an atom of carbon-12 is 12 daltons, which corresponds with the number of nucleons that it has (6 protons and 6 neutrons). However, the mass of an atomic-scale object is affected by the binding energy of the nucleons in its atomic nuclei, as well as the mass and binding energy of its electrons. Therefore, this equality holds only for the carbon-12 atom in the stated conditions, and will vary for other substances. For example, the mass of one unbound atom of the common hydrogen isotope (hydrogen-1, protium) is , the mass of one free neutron is and the mass of one hydrogen-2 (deuterium) atom is . In general, the difference (mass defect) is less than 0.1%; exceptions include hydrogen-1 (about 0.8%), helium-3 (0.5%), lithium (0.25%) and beryllium (0.15%). The unified atomic mass unit and the dalton should not be confused with the unit of mass in the atomic units systems, which is instead the electron rest mass (''m''e).

Energy equivalents

The atomic mass constant can also be expressed as its energy equivalent, that is ''m''u''c''2. The 2018 CODATA recommended values are: : : The megaelectronvolt (MeV) is commonly used as a unit of mass in particle physics, and these values are also important for the practical determination of relative atomic masses.


Origin of the concept

The interpretation of the law of definite proportions in terms of the atomic theory of matter implied that the masses of atoms of various elements had definite ratios that depended on the elements. While the actual masses were unknown, the relative masses could be deduced from that law. In 1803 John Dalton proposed to use the (still unknown) atomic mass of the lightest atom, that of hydrogen, as the natural unit of atomic mass. This was the basis of the atomic weight scale. For technical reasons, in 1898, chemist Wilhelm Ostwald and others proposed to redefine the unit of atomic mass as 1/16 of the mass of an oxygen atom. That proposal was formally adopted by the International Committee on Atomic Weights (ICAW) in 1903. That was approximately the mass of one hydrogen atom, but oxygen was more amenable to experimental determination. This suggestion was made before the discovery of the existence of elemental isotopes, which occurred in 1912. The same definition was adopted in 1909 the physicist Jean Perrin in his extensive experiments to determine the atomic masses and Avogadro's constant. This definition remained unchanged until 1961. Perrin also defined the "mole" as an amount of a compound that contained as many molecules as 32 grams of oxygen (). He called that number the Avogadro number in honor of physicist Amedeo Avogadro.

Isotopic variation

The discovery of isotopes of oxygen in 1929 required a more precise definition of the unit. Unfortunately, two distinct definitions came into use. Chemists choose to define the AMU as 1/16 of the average mass of an oxygen atom as found in nature; that is, the average of the masses of the known isotopes, weighted by their natural abundance. Physicists, on the other hand, defined it as 1/16 of the mass of an atom of the isotope oxygen-16 (16O).

Definition by the IUPAC

The existence of two distinct units with the same name was confusing, and the difference (about in relative terms) was large enough to affect high-precision measurements. Moreover, it was discovered that the isotopes of oxygen had different natural abundances in water and in air. For these and other reasons, in 1961 the International Union of Pure and Applied Chemistry (IUPAC), which had absorbed the ICAW, adopted a new definition of the atomic mass unit for use in both physics and chemistry; namely, 1/12 of the mass of a carbon-12 atom. This new value was intermediate between the two earlier definitions, but closer to the one used by chemists (who would be affected the most by the change). The new unit was named the "unified atomic mass unit" and given a new symbol "u", to replace the old "amu" that had been used for the oxygen-based units. However, the old symbol "amu" has sometimes been used, after 1961, to refer to the new unit, particularly in lay and preparatory contexts. With this new definition, the standard atomic weight of carbon is approximately , and that of oxygen is approximately . These values, generally used in chemistry, are based on averages of many samples from Earth's crust, its atmosphere, and organic materials.

Adoption by the BIPM

The IUPAC 1961 definition of the unified atomic mass unit, with that name and symbol "u", was adopted by the International Bureau for Weights and Measures (BIPM) in 1971 as a non-SI unit accepted for use with the SI.

The dalton

In 1993, the IUPAC proposed the shorter name "dalton" (with symbol "Da") for the unified atomic mass unit. As with other unit names such as watt and newton, "dalton" is not capitalized in English, but its symbol, "Da", is capitalized. The name was endorsed by the International Union of Pure and Applied Physics (IUPAP) in 2005. In 2003 the name was recommended to the BIPM by the Consultative Committee for Units, part of the CIPM, as it "is shorter and works better with he SIprefixes". In 2006, the BIPM included the dalton in its 8th edition of the formal definition of SI. The name was also listed as an alternative to "unified atomic mass unit" by the International Organization for Standardization in 2009. It is now recommended by several scientific publishers, and some of them consider "atomic mass unit" and "amu" deprecated. In 2019, the BIPM retained the dalton in its 9th edition of the formal definition of SI while dropping the unified atomic mass unit from its table of non-SI units accepted for use with the SI, but secondarily notes that the dalton (Da) and the unified atomic mass unit (u) are alternative names (and symbols) for the same unit.

A proposal

In 2012, American engineer Brian Leonard proposed to redefine the dalton (and presumably the unified atomic mass unit) as 1/''N'' grams, thereby breaking the link with 12C. This would imply changes in the atomic masses of all elements when expressed in daltons, but the change would be too small to have practical effects.

2019 redefinition of the SI base units

The definition of the dalton was not affected by the 2019 redefinition of SI base units, that is, 1 Da in the SI is still 1/12 of the mass of a carbon-12 atom, a quantity that must be determined experimentally in terms of SI units. However, the definition of a mole was changed to be the amount of substance consisting of exactly entities and the definition of the kilogram was changed as well. As a consequence, the molar mass constant is no longer exactly 1 g/mol, meaning that the number of grams in the mass of one mole of any substance is no longer exactly equal to the number of daltons in its average molecular mass.


Although relative atomic masses are defined for neutral atoms, they are measured (by mass spectrometry) for ions: hence, the measured values must be corrected for the mass of the electrons that were removed to form the ions, and also for the mass equivalent of the electron binding energy, ''E''b/''m''u''c''2. The total binding energy of the six electrons in a carbon-12 atom is 1030.1089 eV = : ''E''b/''m''u''c''2 = , or about one part in 10 million of the mass of the atom. Before the 2019 redefinition of SI units, experiments were aimed to determine the value of the Avogadro constant for finding the value of the unified atomic mass unit.

Josef Loschmidt

A reasonably accurate value of the atomic mass unit was first obtained indirectly by Josef Loschmidt in 1865, by estimating the number of particles in a given volume of gas.

Jean Perrin

Perrin estimated the Avogadro number by a variety of methods, at the turn of the 20th century. He was awarded the 1926 Nobel Prize in Physics, largely for this work.


The electric charge per mole of electrons is a constant called the Faraday constant, whose value had been essentially known since 1834 when Michael Faraday published his works on electrolysis. In 1910, Robert Millikan obtained the first measurement of the charge on an electron, ''e''. The quotient ''F''/''e'' provided an estimate of Avogadro's number. The classic experiment is that of Bower and Davis at NIST, and relies on dissolving silver metal away from the anode of an electrolysis cell, while passing a constant electric current ''I'' for a known time ''t''. If ''m'' is the mass of silver lost from the anode and ''A'' the atomic weight of silver, then the Faraday constant is given by: :F = \frac. The NIST scientists devised a method to compensate for silver lost from the anode by mechanical causes, and conducted an isotope analysis of the silver used to determine its atomic weight. Their value for the conventional Faraday constant was ''F'' = , which corresponds to a value for the Avogadro constant of : both values have a relative standard uncertainty of .

Electron mass measurement

In practice, the atomic mass constant is determined from the electron rest mass ''m''e and the electron relative atomic mass ''A''r(e) (that is, the mass of electron divided by the atomic mass constant). 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. :m_ = \frac = \frac , :m_ = \frac , :N_ = \frac = \frac where ''c'' is the speed of light, ''h'' is the Planck constant, ''α'' is the fine-structure constant, and ''R'' is the Rydberg constant. As may be observed from the old values (2014 CODATA) in the table below, the main limiting factor in the precision of the Avogadro constant was the uncertainty in the value of the Planck constant, as all the other constants that contribute to the calculation were known more precisely. The power of the presently defined values of universal constants can be understood from the table below(2018 CODATA).

X-ray crystal density methods

thumb|right|200px|Ball-and-stick_model_of_the_[[unit_cell_of_[[silicon.html" style="text-decoration: none;"class="mw-redirect" title="unit_cell.html" style="text-decoration: none;"class="mw-redirect" title="Ball-and-stick model of the [[unit cell">Ball-and-stick model of the [[unit cell of [[silicon">unit_cell.html" style="text-decoration: none;"class="mw-redirect" title="Ball-and-stick model of the [[unit cell">Ball-and-stick model of the [[unit cell of [[silicon. X-ray diffraction measures the cell parameter, ''a'', which is used to calculate a value for the Avogadro constant.]] [[Silicon]] single crystals may be produced today in commercial facilities with extremely high purity and with few lattice defects. This method defined the Avogadro constant as the ratio of the [[molar volume]], ''V'', to the atomic volume ''V'':
:N_ = \frac, where V_ = \frac and ''n'' is the number of atoms per unit cell of volume ''V''cell. The unit cell of silicon has a cubic packing arrangement of 8 atoms, and the unit cell volume may be measured by determining a single unit cell parameter, the length ''a'' of one of the sides of the cube. The 2018 CODATA value of ''a'' for silicon is . In practice, measurements are carried out on a distance known as ''d''(Si), which is the distance between the planes denoted by the Miller index|Miller indices , and is equal to ''a''/. The isotope proportional composition of the sample used must be measured and taken into account. Silicon occurs in three stable isotopes (28Si, 29Si, 30Si), and the natural variation in their proportions is greater than other uncertainties in the measurements. The atomic weight ''A'' for the sample crystal can be calculated, as the standard atomic weights of the three nuclides are known with great accuracy. This, together with the measured density ''ρ'' of the sample, allows the molar volume ''V'' to be determined: :V_ = \frac where ''M'' is the molar mass constant. The 2018 CODATA value for the molar volume of silicon is , with a relative standard uncertainty of .

See also

* Mass (mass spectrometry) ** Kendrick mass ** Monoisotopic mass * Mass-to-charge ratio


Meng Wang, G. Audi, F.G. Kondev, W.J. Huang, S. Naimi, and Xing Xu (2017): "The Ame2016 atomic mass evaluation (II). Tables, graphs and references". ''Chinese Physics C'', volume 41, issue 3, article 030003, pages 1-441. Integrated DNA Technologies (2011):
Molecular Facts and Figures
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IDT website, Support & Education section
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Bureau International des Poids et Mesures (2019):
The International System of Units (SI)
', 9th edition, English version, page 134. Available at th
BIPM website
International Bureau for Weights and Measures (2017):
Proceedings of the 106th meeting of the International Committee for Weights and Measures (CIPM), 16-17 and 20 October 2017
', page 23. Available at th

International Bureau for Weights and Measures (2018):
Resolutions Adopted - 26th Confernce Générale des Poids et Mesures
'. Available at th
BIPM website
Oseen, C.W. (December 10, 1926).
Presentation Speech for the 1926 Nobel Prize in Physics


' From the ''Encyclopaedia Britannica'', 15th edition; reproduced by NIST. Accessed on 2019-07-03.
Bureau International des Poids et Mesures (1971):
14th Conference Générale des Poids et Mesures
' Available at th
BIPM website

External links

Atomic weights and isotopic compositions

at sizes.com {{SI units|state=collapsed Category:Metrology Category:Nuclear chemistry Category:Units of chemical measurement Category:Units of mass