name	unified atomic mass unit
standard	SI recognized unit
quantity	mass
symbol	u
units1	dalton
inunits1	1
units2	kg
inunits2
units3	eV/c2
inunits3
units4	me
inunits4	1822.88839
inunits5	}}

The unified atomic mass unit (Also known as amu, symbol: u) or dalton (symbol: Da) is a unit that is used for indicating mass on an atomic or molecular scale. It is defined as one twelfth of the rest mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state, and has a value of . One dalton is ''approximately'' equal to the mass of one proton or one neutron; an equivalence of saying 1 g mol⁻¹ . The CIPM have categorised it as a "non-SI unit" because units values in SI units must be obtained experimentally.

History[link]

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

The discovery of isotopic oxygen in 1929 led to a divergence in atomic weight representation, with isotopically weighted oxygen (chemistry) and pure ¹⁶O (physics) bases both used as the basis for the atomic mass unit (amu). The inevitable divergence could result in errors in computations, and was thus unwieldy. The reference was changed to carbon-12 in 1961 and a new symbol "u" replaced the now deprecated "amu".

: $1\; \backslash mathrm\{u\}\; =\; m\_\{\backslash mathrm\{u\}\}\; =\; m(\{\}^\{12\}\backslash mathrm\{C\})\; /\; 12\; \backslash ,$

The current unit is referred to as the "unified atomic mass unit" u. The choice of carbon-12 was used to minimise further divergence with prior literature.

Terminology[link]

The unified atomic mass unit and the dalton are different names for the same unit of measure. 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 approved the use of the dalton with the qualification that the GCPM had not given its approval. 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''". 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. 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 units whose values in SI units must be obtained experimentally: Units accepted for use with the SI". 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 ISO 80000-1:2009 (General), identified the dalton as having "''earlier [been] called the unified atomic mass unit u''", but ISO 80000-10:2009 (atomic and nuclear physics) catalogued both as being alternatives for each other. 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.''" 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''".

Relationship to SI[link]

The definition of the mole, an SI base unit, was accepted by the CGPM in 1971 as: # 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". # 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. The definition of the mole also determines the value of the universal constant that relates the number of entities to amount of substance for any sample. This constant is called the Avogadro constant, symbol ''N_A'' or ''L'', and is equal to entities per mole.

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 12u, it follows that there are ''N_A'' atoms of carbon 12 in 0.012 kg of carbon 12. This can be expressed mathematically as :''N_A(12u) = 0.012 kg/mol'', or :''N_Au = 0.001 kg/mol''

Masses of proteins are often expressed in daltons. For example, a protein with a molecular weight of 64000 g mol^-1 has a mass of 64000 daltons or 64 kDa in short.

See also [link]

Kendrick mass

Mass-to-charge ratio

References[link]

External links[link]

atomic mass unit at sizes.com

Category:Metrology Category:Nuclear chemistry Category:Units of chemical measurement Category:Units of mass

ar:وحدة كتل ذرية ast:Unidá de masa atómica be:Атамная адзінка масы bg:Единица за атомна маса bo:རྡུལ་ཕྲན་སྤུས་ཚད་རྩི་གཞི། bs:Jedinica atomske mase br:Unanenn tolz atomek ca:Unitat de massa atòmica cs:Atomová hmotnostní konstanta da:Dalton (kemisk enhed) de:Atomare Masseneinheit et:Aatommassiühik el:Ατομική μονάδα μάζας es:Unidad de masa atómica eo:Atommasa unuo fa:یکای جرم اتمی fr:Unité de masse atomique unifiée fy:Atomêre massa-ienheid ko:원자 질량 단위 hi:परमाण्विक भार इकाई hr:Atomska jedinica mase id:Satuan massa atom it:Unità di massa atomica he:דלטון kk:Массаның атомдық бірлігі lt:Atominės masės vienetas hu:Atomi tömegegység mk:Единица за атомска маса nl:Atomaire massa-eenheid ja:統一原子質量単位 nap:Unità e' massa atomeca no:Atommasseenhet nn:Atommasseeining uz:Massaning atom birligi nds:Atomare Masseneenheit pl:Jednostka masy atomowej pt:Unidade de massa atômica ro:Unitate atomică de masă ru:Атомная единица массы simple:Atomic mass unit sk:Atómová hmotnostná jednotka sl:Enota atomske mase sr:Јединица атомске масе sh:Jedinica atomske mase fi:Atomimassayksikkö sv:Atommassenhet th:หน่วยมวลอะตอม tr:Atomik kütle birimi uk:Атомна одиниця маси ur:جوہری کمیتی اکائی vi:Đơn vị khối lượng nguyên tử zh:原子质量单位

The atomic mass (m_a) is the mass of a specific isotope, most often expressed in unified atomic mass units. The atomic mass is the total mass of protons, neutrons and electrons in a single atom.

The atomic mass is sometimes incorrectly used as a synonym of relative atomic mass, average atomic mass and atomic weight; these differ subtly from the atomic mass. The atomic mass is defined as the mass of an atom, which can only be one isotope at a time and is not an abundance-weighted average as in the case of atomic weight. In the case of many elements that have one dominant isotope the actual numerical similarity/difference between the atomic mass of the most common isotope and the relative atomic mass or standard atomic weights can be very small such that it does not affect most bulk calculations—but such an error can be critical when considering individual atoms. For elements with more than one common isotope the difference even to the most common atomic mass can be half a mass unit or more (e.g. chlorine). The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.

Standard atomic weight refers to the mean relative atomic mass of an element in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances. These are what are included in a standard periodic table and is what is used in most bulk calculations. An uncertainty in brackets is included which often reflects natural variability in isotopic distribution rather than uncertainty in measurement. For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning. Therefore, for synthetic elements the total nucleon count of the most stable isotope (i.e., the isotope with the longest half-life) is listed in brackets in place of the standard atomic weight. Lithium represents a unique case where the natural abundances of the isotopes have been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources, such as rivers.

Relative atomic mass is a synonym for atomic weight and closely related to average atomic mass (but not a synonym for atomic mass), the weighted mean of the atomic masses of all the atoms of a chemical element found in a particular sample, weighted by isotopic abundance. This is frequently used as a synonym for the standard atomic weight and it is correct to do so since the standard atomic weights are relative atomic masses, although it is less specific to do so. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the average or have different certainties (number of significant figures) than the standard atomic weights.

Relative isotopic mass is the relative mass of a given isotope (more specifically, any single nuclide), scaled with carbon-12 as exactly 12. No other nuclides other than carbon-12 have exactly whole-number masses in this scale. This is due to two factors: [1] the different mass of neutrons and protons acting to change the total mass in nuclides with proton/neutron ratios other than the 1:1 ratio of carbon-12; and [2] an exact whole-number will not be located if there exists a loss/gain of mass to difference in mean binding energy relative to the mean binding energy for carbon-12. Any mass defect due to binding energy is a small fraction (less than 1%) compared to the mass of a nucleon, and even less compared to the average mass per nucleon in carbon-12, which is moderately strongly bound. Since protons and neutrons differ in mass from each by an even smaller fraction (about 0.0014 u), the practice of rounding the atomic mass of any given nuclide or isotope to the nearest whole number, always gives the simple whole number total nucleon count. Neutron count can then be derived by subtracting the atomic number.

The mass number of a nuclide is simply the total number of nucleons in the nucleus. It is equal to the number of protons (atomic number) plus the number of neutrons. This number is always a simple whole number. It has units of "nucleons" not atomic mass units. An example is oxygen-16, which has 16 nucleons (8 protons and 8 neutrons).

Mass defects in atomic masses[link]

The amount that the atomic masses deviate from their mass numbers is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the following: nuclear fission in an element heavier than iron produces energy, and fission in any element lighter than iron requires energy. The opposite is true of nuclear fusion reactions: fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

Measurement of atomic masses[link]

Direct comparison and measurement of the masses of atoms is achieved with mass spectrometry.

Conversion factor between atomic mass units and grams[link]

The standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are in 12 grams of the carbon isotope C-12. The number of atoms in a mole is called Avogadro's number, the value of which is approximately 6.022 × 10 mol⁻¹. One mole of a substance always contains almost exactly the ''relative atomic mass'' or ''molar mass'' of that substance (which is the concept of molar mass), expressed in grams; however, this is almost never true for the ''atomic mass''. For example, the standard atomic weight of iron is 55.847 g/mol, and therefore one mole of iron as commonly found on earth has a mass of 55.847 grams. The ''atomic mass'' of an ⁵⁶Fe isotope is 55.935 u and one mole of ⁵⁶Fe will in theory weigh 55.935g, but such amounts of pure ⁵⁶Fe have never been found on Earth.

The formulaic conversion between atomic mass units and SI mass in grams for a single atom is:

::$1\backslash \; \{\backslash rm\{u\}\}=\{M\_\{\backslash rm\{u\}\}\; \backslash over\; N\_\{\backslash rm\; A\}\}\backslash \; =\backslash over\; N\_\{\backslash rm\; A\}\}$

where $M\_\{\backslash rm\; u\}$ is the Molar mass constant and $N\_\{\backslash rm\; A\}$ is the Avogadro constant.

Relationship between atomic and molecular masses[link]

Similar definitions apply to molecules. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms (nuclides). One can compute the molar mass of a compound by adding the relative atomic masses of the elements given in the chemical formula. In both cases the multiplicity of the atoms (the number of times it occurs) must be taken into account, usually by multiplication of each unique mass by its multiplicity.

History[link]

The first scientists to determine atomic weights were John Dalton and Thomas Thomson between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Atomic weight was originally defined relative to that of the lightest element hydrogen taken as 1.00, and in the 1820s Prout's hypothesis stated that atomic masses of all elements would prove via a whole number rule to be exact multiples of this hydrogen weight. Berzelius, however, soon proved that this hypothesis did not always hold even approximately, and in some elements, such as chlorine, atomic weight falls almost exactly between two multiples of the hydrogen weight. Still later, as noted, this was shown to be an isotope effect, and that the atomic masses of pure isotopes, or nuclides, are multiples of the hydrogen mass, to within about 1%.

In the 1860s Stanislao Cannizzaro refined atomic weights by applying Avogadro's law (notably at the Karlsruhe Congress of 1860). He formulated a law to determine atomic weights of elements: ''the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight'' and determined atomic weights and molecular weights by comparing the vapor density of a collection of gases with molecules containing one or more of the chemical element in question.

In the early twentieth century, up until the 1960s chemists and physicists used two different atomic mass scales. The chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to 2 different tables of atomic mass. The unified scale based on carbon-12, ¹²C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the chemists' scale.

The term ''atomic weight'' is being phased out slowly and being replaced by relative atomic mass, in most current usage. The history of this shift in nomenclature reaches back to the 1960s and has been the source of much debate in the scientific community. The debate was largely created by the adoption of the unified atomic mass unit and the realization that weight was in some ways an inappropriate term. The argument for keeping the term "atomic weight" was primarily that it was a well understood term to those in the field, that the term "atomic mass" was already in use (as it is currently defined) and that the term "relative atomic mass" was in some ways redundant. In 1979, in a compromise move, the definition was refined and the term "relative atomic mass" was introduced as a secondary synonym. Twenty years later the primacy of these synonyms was reversed and the term "relative atomic mass" is now the preferred term; however the "standard atomic weights" have maintained the same name.

See also[link]

Atomic number

Atomic mass unit

Isotope

Isotope geochemistry

Molecular mass

Jean Stas

References[link]

External links[link]

NIST relative atomic masses of all isotopes and the standard atomic weights of the elements

Tutorial on the concept and measurement of atomic mass

AME2003 Atomic Mass Evaluation from the National Nuclear Data Center

Mass, atomic Category:Chemical properties Category:Mass Category:Stoichiometry

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