          # Atomic mass

• this article needs to be updated. in particular: it needs to reflect the 2019 redefinition of the si base units, which came into effect on 20 may 2019. please update this article to reflect recent events or newly available information. (january 2020) stylized lithium-7 atom: 3 protons, 4 neutrons, and 3 electrons (total electrons are ~​14300th 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.

the atomic mass (ma or m) is the mass of an atom. although the si unit of mass is kilogram (symbol: kg), the atomic mass is often expressed in the non-si unit dalton (symbol: da, or u) where 1 dalton is defined as ​112 of the mass of a single carbon-12 atom, at rest. the protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. thus, the atomic mass measured in da has nearly the same value as the mass number. conversion between mass in kg and mass in da can be done using the atomic mass constant .

the formula used for conversion is: where is the molar mass constant, is the avogadro constant and is the experimentally determined molar mass of carbon-12.

the relative isotopic mass (see section below) can be obtained by dividing the atomic mass ma of an isotope by the atomic mass constant mu yielding a dimensionless value. thus, the atomic mass of a carbon-12 atom is 12 da (or 12 u), but the relative isotopic mass of a carbon-12 atom is simply 12. the sum of relative isotopic masses of all atoms in a molecule is the relative molecular mass.

the atomic mass of an isotope and the relative isotopic mass refers to a certain specific isotope of an element. because usually substances are not isotopially pure, it is convenient to use the elemental atomic mass which is the average (mean) atomic mass of an element, weighted by the abundance of the isotopes. the dimensionless (standard) atomic weight is the weighted mean relative isotopic mass of a (typical naturally-occurring) mixture of isotopes.

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 (as per e = mc2).

• relative isotopic mass
• similar terms for different quantities
• mass defects in atomic masses
• measurement of atomic masses
• relationship between atomic and molecular masses
• history
## Not to be confused with Standard atomic weight, Mass number, or Relative atomic mass. For the band, see Atomic Mass (band). This article needs to be updated. In particular: it needs to reflect the 2019 redefinition of the SI base units, which came into effect on 20 May 2019. Please update this article to reflect recent events or newly available information. (January 2020) Stylized lithium-7 atom: 3 protons, 4 neutrons, and 3 electrons (total electrons are ~​1⁄4300th 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. The atomic mass (ma or m) is the mass of an atom. Although the SI unit of mass is kilogram (symbol: kg), the atomic mass is often expressed in the non-SI unit dalton (symbol: Da, or u) where 1 dalton is defined as ​1⁄12 of the mass of a single carbon-12 atom, at rest. The protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. Thus, the atomic mass measured in Da has nearly the same value as the mass number. Conversion between mass in kg and mass in Da can be done using the atomic mass constant $m_{\rm {u}}={{m({\rm {^{12}C}})} \over {12}}=1\ {\rm {Da}}$ . The formula used for conversion is: $1\ {\rm {Da}}=m_{\rm {u}}={M_{\rm {u}} \over {N_{\rm {A}}}}={M(^{12}C) \over {12\ N_{\rm {A}}}}=1.66053906660(50)\cdot 10^{-27}\ {\rm {kg}}$ where $M_{\rm {u}}$ is the molar mass constant, $N_{\rm {A}}$ is the Avogadro constant and $M(^{12}C)$ is the experimentally determined molar mass of carbon-12. The relative isotopic mass (see section below) can be obtained by dividing the atomic mass ma of an isotope by the atomic mass constant mu yielding a dimensionless value. Thus, the atomic mass of a carbon-12 atom is 12 Da (or 12 u), but the relative isotopic mass of a carbon-12 atom is simply 12. The sum of relative isotopic masses of all atoms in a molecule is the relative molecular mass. The atomic mass of an isotope and the relative isotopic mass refers to a certain specific isotope of an element. Because usually substances are not isotopially pure, it is convenient to use the elemental atomic mass which is the average (mean) atomic mass of an element, weighted by the abundance of the isotopes. The dimensionless (standard) atomic weight is the weighted mean relative isotopic mass of a (typical naturally-occurring) mixture of isotopes. 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 (as per E = mc2). Contents 1 Relative isotopic mass 2 Similar terms for different quantities 3 Mass defects in atomic masses 4 Measurement of atomic masses 5 Relationship between atomic and molecular masses 6 History 7 See also 8 References 9 External links  