Instruct: Given a question, retrieve passages that can help answer the question.
Query: What is 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.[1] 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.
Whereas the mass number simply counts the total number of neutrons and protons and is thus a natural (or whole) number, the atomic mass of a single atom is a real number giving the mass of a particular isotope (or "nuclide") of the element, expressed in atomic mass units (symbol: u). In general, the mass number of a given nuclide differs in value slightly from its atomic mass, since the mass of each proton and neutron is not exactly 1 u; since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number; and (finally) because of the nuclear binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969 u and that of chlorine-37 is 36.966 u. However, the atomic mass in u of each isotope is quite close to its simple mass number (always within 1%). The only isotope whose atomic mass is exactly a natural number is C, which by definition has a mass of exactly 12 because u is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.
The atomic mass constant can also be expressed as its energy equivalent, that is "m""c". The 2010 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. Although relative atomic masses are defined for neutral atoms, they are measured (by mass spectrometry) for ions: hence, the measured values must be correct 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"/"m""c". The total binding energy of the six electrons in a carbon-12 atom is 1030.1089 eV = : "E"/"m""c" = , or about one part in 10 million of the mass of the atom.
Atomic mass ("m") is the mass of a single atom, with unit "Da" or "u" (the unified atomic mass unit). It defines the mass of a specific isotope, which is an input value for the determination of the relative atomic mass. An example for three silicon isotopes is given below.
The atomic mass and relative isotopic mass are sometimes confused, or incorrectly used, as synonyms of standard atomic weight (also known as atomic weight) and the standard atomic weight (a particular variety of atomic weight, in the sense that is a standardized atomic weight). However, as noted in the introduction, atomic weight and standard atomic weight represent terms for (abundance-weighted) averages of atomic masses in elemental samples, not for single nuclides. As such, atomic weight and standard atomic weight often differ numerically from relative isotopic mass and atomic mass, and they can also have different units than atomic mass when this quantity is not expressed in unified atomic mass units (see the linked article for atomic weight).
It is a synonym for atomic weight, though it is not to be confused with relative isotopic mass. Relative atomic mass is also frequently used as a synonym for "standard" atomic weight and these quantities may have overlapping values if the relative atomic mass used is that for an element from Earth under defined conditions. However, relative atomic mass (atomic weight) is still technically distinct from standard atomic weight because of its application only to the atoms obtained from a single sample; it is also not restricted to terrestrial samples, whereas standard atomic weight averages multiple samples but only from terrestrial sources. Relative atomic mass is therefore a more general term that can more broadly refer to samples taken from non-terrestrial environments or highly specific terrestrial environments which may differ substantially from Earth-average or reflect different degrees of certainty (e.g. in number of significant figures) than those reflected in standard atomic weights.
One mole of a substance always contains almost exactly the standard atomic weight 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 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 55.847grams. The atomic mass of the 56Fe isotope is 55.935 u and one mole of 56Fe atoms would then in theory have a mass of 55.935 g, but such amounts of pure 56Fe have never been found (or separated out) on Earth. However, there are 22 mononuclidic elements of which essentially only a single isotope is found in nature (common examples are fluorine, sodium, aluminum and phosphorus) and for these elements the standard atomic weight and atomic mass are the same. Samples of these elements therefore may serve as reference standards for certain atomic mass values.
In practice, the atomic mass constant is determined as the ratio of the electron rest mass "m" to the electron relative atomic mass "A"(e) (that is, the mass of the electron on a scale where C = 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.
The current uncertainty in the value of the atomic mass constant – one part in 20 million – is almost entirely due to the uncertainty in the value of the Planck constant.
The electron relative atomic mass also enters into the calculation of all other relative atomic masses. By convention, relative atomic masses are quoted for neutral atoms, but the actual measurements are made on positive ions, either in a mass spectrometer or a Penning trap. Hence the mass of the electrons must be added back on to the measured values before tabulation. A correction must also be made for the mass equivalent of the binding energy "E". Taking the simplest case of complete ionization of all electrons, for a nuclide X of atomic number "Z",
As relative atomic masses are measured as ratios of masses, the corrections must be applied to both ions: the uncertainties in the corrections are negligible, as illustrated below for hydrogen 1 and oxygen 16.
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 12 daltons or 12 unified atomic mass units. 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 is about 1.998467052 × 10−26 kilograms.
Relative atomic mass is determined by the "average atomic mass", or the weighted mean of the atomic masses of all the atoms of a particular chemical element found in a particular sample, which is then compared to the atomic mass of carbon-12. This comparison is the quotient of the two weights, which makes the value dimensionless (no unit appended). This quotient also explains the word "relative": the sample mass value is considered relative to that of carbon-12..
The sum of the atomic number "Z" and the number of neutrons, "N", gives the mass number "A" of an atom. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes) and the mass defect of nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in unified atomic mass units (making a quantity called the "relative isotopic mass"), is within 1% of the whole number "A".
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 actual mass of an atom at rest is often expressed using the unified atomic mass unit (u), also called dalton (Da). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately . Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 u. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 u). However, this number will not be exactly an integer except in the case of carbon-12 (see below). The heaviest stable atom is lead-208, with a mass of .
The mass number should also not be confused with the standard atomic weight (also called atomic weight) of an element, which is the ratio of the average atomic mass of the different isotopes of that element (weighted by abundance) to the unified atomic mass unit. The atomic weight is an actual "mass" (made relative, i.e., a ratio), while the mass number is a "counted" number (and so an integer). 
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 (as it is currently defined) and that the term "relative atomic mass" might be easily confused with relative 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).
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 1g/mol. 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, and has a value of , or approximately 1.66 yoctograms. 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.
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. The only difference in that case, is that 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.
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 of chlorine 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 atomic mass (symbol: "A") 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 one unified atomic mass unit. The unified atomic mass unit (symbol: "u" or "Da") is defined as being of the atomic mass of a carbon-12 atom. Since both values in the ratio are expressed in the same unit ("u"), the resulting value is dimensionless; hence the value is said to be "relative".
The unified atomic mass unit (symbol: u) is the standard unit that is used for indicating mass on an atomic or molecular scale (atomic mass). The dalton (symbol: Da) is equivalent to the unified atomic mass unit. One unified atomic mass unit is approximately the mass of one a single proton or neutron. The unified atomic mass unit has a value of . The "amu" without the "unified" prefix is an obsolete unit based on oxygen, which was replaced in 1961.
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 H. Despite the initial mass of H 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 mass of oxygen (). This evaluation was made prior to the discovery of the existence of elemental isotopes, which occurred in 1912.
Any mass defect due to nuclear binding energy is experimentally a small fraction (less than 1%) of the mass of 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 (about 0.0014 u), 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).
Although all elements are the product of nuclear fusion of hydrogen into higher elements, it is now understood that atoms consist of both protons (hydrogen nuclei) and neutrons. The modern version of Prout's rule is that the atomic mass of an isotope of proton number (atomic number) "Z" and neutron number "N" is equal to sum of the masses of its constituent protons and neutrons, minus the mass of the nuclear binding energy, the mass defect. According to the whole number rule proposed by Francis Aston, the mass of an isotope is roughly, but not exactly, its mass number "A" = "Z" + "N" times an atomic mass unit (u), plus or minus binding energy discrepancy – atomic mass unit being the modern approximation for "mass of a proton, neutron, or hydrogen atom". For example iron-56 atoms (which have among the highest binding-energies) weigh only about 99.1% as much as 56 hydrogen atoms. The missing 0.9% of mass represents the energy lost when the nucleus of iron was made from hydrogen inside a star. (See stellar nucleosynthesis).
Both atomic and molecular masses are usually obtained relative to the mass of the isotope C (carbon 12), which by definition is equal to 12. For example, the molecular mass of methane, whose molecular formula is CH, is calculated as follows:
A more proper term would be "relative molecular mass". However the adjective 'relative' is omitted as it is universally assumed that atomic and molecular masses are relative to the mass of C. Relative atomic and molecular mass values are dimensionless but are given the "unit" Dalton (formerly atomic mass unit) to indicate that the number is equal to the mass of one molecule divided by of the mass of one atom of C. The mass of 1 mol of substance is designated as molar mass. By definition, it has the unit gram.