Editing Chemical element (section)

=== Isotopic mass and atomic mass ===
{{main|atomic mass|relative atomic mass}}

The [[mass number]] of an element, ''A'', is the number of [[nucleon]]s (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass number, which is written as a superscript on the left hand side of the chemical symbol (e.g., {{sup|238}}U). The mass number is always an integer and has units of "nucleons". Thus, [[Isotopes of magnesium|magnesium-24]] (24 is the mass number) is an atom with 24 nucleons (12 protons and 12 neutrons).

Whereas the mass number simply counts the total number of neutrons and protons and is thus an integer, the [[atomic mass]] of a particular isotope (or "nuclide") of the element is the mass of a single atom of that isotope, and is typically expressed in [[Dalton (unit)|dalton]]s (symbol: Da), aka universal atomic mass units (symbol: u). Its [[relative atomic mass]] is a dimensionless number equal to the atomic mass divided by the [[atomic mass constant]], which equals 1&nbsp;Da. In general, the mass number of a given nuclide differs in value slightly from its relative atomic mass, since the mass of each proton and neutron is not exactly 1&nbsp;Da; since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number; and because of the [[nuclear binding energy]] and electron binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969&nbsp;Da and that of chlorine-37 is 36.966&nbsp;Da. However, the relative atomic mass of each isotope is quite close to its mass number (always within 1%). The only isotope whose atomic mass is exactly a [[natural number]] is {{sup|12}}C, which has a mass of 12&nbsp;Da; because the dalton is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.

The [[standard atomic weight]] (commonly called "atomic weight") of an element is the ''average'' of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit. This number may be a fraction that is ''not'' close to a whole number. For example, the relative atomic mass of chlorine is 35.453&nbsp;u, which differs greatly from a whole number as it is an average of about 76% chlorine-35 and 24% chlorine-37. Whenever a relative atomic mass value differs by more than ~1% from a whole number, it is due to this averaging effect, as significant amounts of more than one isotope are naturally present in a sample of that element.