Stylizedlithium-7 atom: 3 protons, 4 neutrons, and 3 electrons (total electrons are ~1/4300 of the mass of the nucleus). It has a mass of 7.016 Da. Rare lithium-6 (mass of 6.015 Da) has only 3 neutrons, reducing the atomic weight (average) of lithium to 6.941.
Atomic mass is often measured indalton (Da) orunified atomic mass unit (u). One dalton is equal to+1/12 the mass of acarbon-12 atom in its natural state, given by theatomic mass constantmu =m(12C)/12 = 1 Da, wherem(12C) is the atomic mass of carbon-12. Thus, the numerical value of the atomic mass of anuclide when expressed in daltons is close to itsmass number.
Therelative isotopic mass (see section below) can be obtained by dividing the atomic massma of anisotope by the atomic mass constantmu, yielding adimensionless value. Thus, the atomic mass of a carbon-12 atomm(12C) is12 Da by definition, but the relative isotopic mass of a carbon-12 atomAr(12C) is simply 12. The sum ofrelative isotopic masses of all atoms in a molecule is therelative molecular mass.
The atomic mass of an isotope and therelative isotopic mass refers to a certain specific isotope of an element. Because substances are usually not isotopically pure, it is convenient to use theelemental atomic mass which is theaverage atomic mass of an element, weighted by the abundance of the isotopes. The dimensionless(standard) atomic weight is theweighted mean relative isotopic mass of a (typical naturally occurring) mixture of isotopes.
Relativeisotopic mass (a property of a single atom) is not to be confused with the averaged quantityatomic 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 scalingrelative to carbon-12.
The relative isotopic mass, then, is the mass of a given isotope (specifically, any singlenuclide), when this value is scaled by the mass ofcarbon-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 12daltons. 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 is1.99264688270(62)×10−26 kg.
As is the case for the relatedatomic mass when expressed indaltons, 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.
The atomic mass or relative isotopic mass are sometimes confused, or incorrectly used, as synonyms ofrelative atomic mass (also known as atomic weight) or thestandard 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. As such, relative atomic mass and standard atomic weight often differ numerically from the relative isotopic mass.
The atomic mass (relative isotopic mass) is defined as the mass of a single atom, which can only be oneisotope (nuclide) at a time, 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 in principle be measured to high precision, since every specimen of such a nuclide is expected to be exactly identical to every other specimen, as all atoms of a given type in the same energy state, and every specimen of a particular nuclide, are expected to be exactly identical in mass to every other specimen of that nuclide. For example, every atom of oxygen-16 is expected to have exactly the same atomic mass (relative isotopic mass) as every other atom of oxygen-16.
In the case of many 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 can be small or even nil, and does not affect most bulk calculations. However, such an error can exist and even be important when considering individual atoms for elements that are not mononuclidic.
For non-mononuclidic 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 ofchlorine 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:
protons and neutrons have different masses,[2][3] and different nuclides have different ratios of protons and neutrons.
atomic masses are reduced, to different extents, by theirbinding energies.
The ratio of atomic mass tomass number (number of nucleons) varies from0.9988381346(51) for56Fe to1.007825031898(14) for1H.
Anymass defect due tonuclear binding energy is experimentally a small fraction (less than 1%) of the mass of an equal number of free nucleons. When compared to the average mass per nucleon in carbon-12, which is moderately strongly bound compared with other atoms, the mass defect of binding for most atoms is an even smaller fraction of a dalton (unified atomic mass unit, based on carbon-12). Since free protons and neutrons differ from each other in mass by a small fraction of a dalton (1.38844933(49)×10−3 Da),[4] rounding the relative isotopic mass, or the atomic mass of any given nuclide given in daltons to the nearest whole number, always gives the nucleon count, or mass number. Additionally, the neutron count (neutron number) may then be derived by subtracting the number of protons (atomic number) from the mass number (nucleon count).
Binding energy per nucleon of common isotopes. A graph of the ratio of mass number to atomic mass would be similar.
The amount that the ratio of atomic masses to mass number deviates from 1 is as follows: the deviation starts positive athydrogen-1, then decreases until it reaches a local minimum at helium-4. Isotopes of lithium, beryllium, and boron are less strongly bound than helium, as shown by their increasing mass-to-mass number ratios.
At carbon, the ratio of mass (in daltons) to mass number is defined as 1, and after carbon it becomes less than one until a minimum is reached atiron-56 (with only slightly higher values for iron-58 andnickel-62), then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the fact thatnuclear fission in an element heavier thanzirconium produces energy, and fission in any element lighter thanniobium requires energy. On the other hand,nuclear fusion of two atoms of an element lighter thanscandium (except for helium) produces energy, whereas fusion in elements heavier thancalcium requires energy. The fusion of two atoms of4He yieldingberyllium-8 would require energy, and the beryllium would quickly fall apart again.4He can fuse withtritium (3H) or with3He; these processes occurred duringBig Bang nucleosynthesis. The formation of elements with more than seven nucleons requires the fusion of three atoms of4He in thetriple-alpha process, skipping over lithium, beryllium, and boron to produce carbon-12.
Here are some values of the ratio of atomic mass to mass number:[5]
Similar definitions apply tomolecules. One can calculate themolecular mass of a compound by adding the atomic masses (not the standard atomic weights) of its constituent atoms. Conversely, themolar mass is usually computed from thestandard atomic weights (not the atomic or nuclide masses). Thus, molecular mass and molar mass differ slightly in numerical value and represent different concepts. Molecular mass is the mass of a molecule, which is the sum of its constituent atomic masses. Molar mass is an average of the masses of the constituent molecules in a chemically pure but isotopically heterogeneous ensemble. 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.
Molar mass of CH4
Standard atomic weight
Number
Total molar mass (g/mol) or molecular weight (unitless)
The first scientists to determine relative atomic masses wereJohn Dalton andThomas Thomson between 1803 and 1805 andJöns Jakob Berzelius between 1808 and 1826. Relative atomic mass (Atomic weight) was originally defined relative to that of the lightest element, hydrogen, which was taken as 1.00, and in the 1820s,Prout's hypothesis stated that atomic masses of all elements would prove to be exact multiples of that of hydrogen. Berzelius, however, soon proved that this was not even approximately true, and for some elements, such as chlorine, relative atomic mass, at about 35.5, falls almost exactly halfway between two integral multiples of that of hydrogen. Still later, this was shown to be largely due to a mix of isotopes, and that the atomic masses of pure isotopes, ornuclides, are multiples of the hydrogen mass, to within about 1%.
In the 1860s,Stanislao Cannizzaro refined relative atomic masses by applyingAvogadro's law (notably at theKarlsruhe Congress of 1860). He formulated a law to determine relative atomic masses of elements:the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight and determined relative atomic masses and molecular masses by comparing thevapor density of a collection of gases with molecules containing one or more of the chemical element in question.[6]
In the 20th century, until the 1960s, chemists and physicists used two different atomic-mass scales. The chemists used an "atomic mass unit" (amu) scale such that the natural mixture ofoxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to only the atomic mass of the most common oxygen isotope (16O, containing eight protons and eight neutrons). However, becauseoxygen-17 andoxygen-18 are also present in natural oxygen this led to two 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. This was adopted as the 'unified atomic mass unit'. The current International System of Units (SI) primary recommendation for the name of this unit is thedalton and symbol 'Da'. The name 'unified atomic mass unit' and symbol 'u' are recognized names and symbols for the same unit.[7]
The termatomic weight is being phased out slowly and being replaced byrelative atomic mass, in most current usage. This shift in nomenclature reaches back to the 1960s and has been the source of much debate in the scientific community, which was triggered by the adoption of theunified 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" might be easily confused withrelative isotopic mass (the mass of a single atom of a given nuclide, expressed dimensionlessly relative to 1/12 of the mass of carbon-12; see section above).
In 1979, as a compromise, the term "relative atomic mass" was introduced as a secondary synonym for atomic weight. Twenty years later the primacy of these synonyms was reversed, and the term "relative atomic mass" is now the preferred term.
However, the term "standard atomic weights" (referring to the standardized expectation atomic weights of differing samples) has not been changed,[8] because simple replacement of "atomic weight" with "relative atomic mass" would have resulted in the term "standard relative atomic mass".
^"Proton mass in u".The NIST Reference on Constants, Units, and Uncertainty. May 2019.Archived from the original on 2000-12-07. Retrieved24 June 2021.
^"neutron mass in u".The NIST Reference on Constants, Units, and Uncertainty. May 2019.Archived from the original on 2000-12-07. Retrieved24 June 2021.
