The atomic mass is the mass of an atomic particle, sub-atomic particle, or molecule. It may be
expressed in unified atomic mass units; by international agreement, 1 atomic mass unit
is defined as 1/12 of the mass of a single carbon-12 atom. When expressed in such units,
the atomic mass is called the relative isotopic mass.
The atomic mass or relative isotopic mass refers to the mass of a single particle, and
is fundamentally different from the quantities elemental atomic weight and standard atomic
weight, both of which refer to averages of naturally-occurring atomic mass values for
samples of elements. Most elements have more than one stable nuclide; for those elements,
such an average depends on the mix of nuclides present, which may vary to some limited extent
depending on the source of the sample, as each nuclide has a different mass. By contrast,
atomic mass figures refer to an individual particle species: as atoms of the same species
are identical, atomic mass values are expected to have no intrinsic variance at all. Atomic
mass figures are thus commonly reported to many more significant figures than atomic
weights. 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.
Relative isotopic mass: the same quantity as atomic mass, but with different units
Relative isotopic mass is not to be confused with the averaged quantity "relative atomic
mass," which is the same as atomic weight. Relative isotopic mass is similar to atomic
mass and has exactly the same numerical value as atomic mass, whenever atomic mass is expressed
in unified atomic mass units. However, relative isotopic mass is a pure 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, when this value is scaled by the mass of carbon-12, when the latter
is set equal to 12. 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 carbon-12 is exactly 12, but no nuclides other
than carbon-12 have exactly whole-number values in this scale. As is the case for the related
atomic mass when expressed in unified atomic mass units, 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 more fully below. Similar terms for different quantities
The atomic mass and relative isotopic mass are sometimes confused, or incorrectly used,
as synonyms of relative atomic mass and the standard atomic weight. However, as noted
in the introduction, atomic weight and standard atomic weight represent terms for averages
of isotopic abundances, not for single nuclides. As such, atomic weight and standard atomic
weight often differ numerically from relative isotopic mass, and they can also have different
units than atomic mass when this quantity is not expressed in unified atomic mass units.
The atomic mass is defined as the mass of a single atom, which can only be one isotope
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 a very
great 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 as every other atom of oxygen-16.
In the case of many elements that have one naturally occurring isotope or one dominant
isotope, the actual numerical similarity/difference between the atomic mass of the most common
isotope, and the relative atomic mass or atomic weight can be small or even nil, and does
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 from even the most common relative isotopic mass, can be half
a mass unit or more. The atomic mass of an uncommon isotope can differ from the relative
atomic mass, atomic weight, or standard atomic weight, by several mass units.
Atomic masses expressed in unified atomic mass units are always close to whole-number
values, but never exactly a whole number, for two reasons:
protons and neutrons have different masses, and different nuclides have different ratios
of protons and neutrons. atomic masses are reduced, to different extents,
by their binding energies. The ratio of atomic mass to mass number varies
from about 0.99884 for 56Fe to 1.00782505 for 1H.
Any mass defect due to nuclear binding energy is experimentally a small fraction of the
mass of a nucleon, and is an even smaller percentage compared to the average mass per
nucleon in carbon-12, which is moderately strongly-bound compared with other atoms.
Since protons and neutrons differ from each other in mass by an even smaller fraction,
rounding the atomic mass of any given nuclide to the nearest whole number always gives the
nucleon count, or mass number. The neutron count may then be derived by subtracting the
number of protons. Mass defects in atomic masses
The amount that the ratio of atomic masses to mass number deviates from 1 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 fact that nuclear fission in an element
heavier than zirconium produces energy, and fission in any element lighter than niobium
requires energy. On the other hand, nuclear fusion reactions: fusion of two atoms of an
element lighter than scandium produces energy, whereas fusion in elements heavier than calcium
requires energy. The difference mass number minus atomic mass
is not maximized at iron-56, but at heavier elements where it reaches about 0.10 atomic
mass units. Here are some values of the ratio of atomic
mass to mass number: Measurement of atomic masses
Direct comparison and measurement of the masses of atoms is achieved with mass spectrometry.
Conversion factor between atomic mass units and grams
The standard scientific unit used to quantify the amount of a substance in macroscopic quantities
is the mole, which is defined arbitrarily as the amount of a substance which has as
many atoms or molecules as there are atoms 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 × 1023. One mole of a substance always contains almost
exactly the relative atomic mass or molar mass of that substance; however, this may
or may not be true for the atomic mass, depending on whether or not the element exists naturally
in more than one isotope. For example, the relative atomic mass 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 the 56Fe isotope is 55.935 u and one mole of 56Fe atoms would then in
theory weigh 55.935 g, but such amounts of pure 56Fe have never been found on Earth.
However there are 22 mononuclidic elements of which essentially only a single isotope
is found in nature and for these elements the relative atomic mass and atomic mass are
the same. Samples of these elements therefore may serve as reference standards for certain
atomic mass values. The formula for conversion between atomic
mass units and SI mass in grams for a single atom is:
where is the Molar mass constant and is the Avogadro constant.
Relationship between atomic and molecular masses
Similar definitions apply to molecules. One can compute the molecular mass of a compound
by adding the atomic masses of its constituent atoms. 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 must be taken into account, usually by multiplication
of each unique mass by its multiplicity. History
The first scientists to determine relative atomic masses were John Dalton and Thomas
Thomson between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Relative
atomic mass 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, or nuclides, are multiples
of the hydrogen mass, to within about 1%. In the 1860s Stanislao Cannizzaro refined
relative atomic masses by applying Avogadro's law. 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 the vapor density of a collection of gases with molecules
containing one or more of the chemical element in question.
In the 20th century, 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. However, because oxygen-17 and oxygen-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. The term atomic weight is being phased out
slowly and being replaced by relative 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 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 and that the term "relative
atomic mass" was in some ways redundant. In 1979, as a compromise, 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
Atomic number Atomic mass unit
Isotope Isotope geochemistry
Molecular mass Jean Stas
References
External links NIST relative atomic masses of all isotopes
and the standard atomic weights of the elements AME2003 Atomic Mass Evaluation from the National
Nuclear Data Center