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Relative atomic mass (symbol: A r; sometimes abbreviated RAM or r.a.m.), also known by the deprecated synonym 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 the atomic mass constant.
Thus, the atomic mass of a carbon-12 atom is 12 Da by definition, 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.
The mass number gives an estimate of the isotopic mass measured in atomic mass units (u). For 12 C, the isotopic mass is exactly 12, since the atomic mass unit is defined as 1/12 of the mass of 12 C. For other isotopes, the isotopic mass is usually within 0.1 u of the mass number. For example, 35 Cl (17 protons and 18 neutrons) has a mass ...
The standard atomic weight of a chemical element (symbol A r °(E) for element "E") is the weighted arithmetic mean of the relative isotopic masses of all isotopes of that element weighted by each isotope's abundance on Earth. For example, isotope 63 Cu (A r = 62.929) constitutes 69% of the copper on Earth, the rest being 65 Cu (A r = 64.927), so
The exact mass of an isotopic species (more appropriately, the calculated exact mass [9]) is obtained by summing the masses of the individual isotopes of the molecule. For example, the exact mass of water containing two hydrogen-1 (1 H) and one oxygen-16 (16 O) is 1.0078 + 1.0078 + 15.9949 = 18.0105 Da.
Silicon (14 Si) has 25 known isotopes, with mass numbers ranging from 22 to 46. 28 Si (the most abundant isotope, at 92.23%), 29 Si (4.67%), and 30 Si (3.1%) are stable. The longest-lived radioisotope is 32 Si, which is produced by cosmic ray spallation of argon.
The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet, and even from place to place on the Earth, but remains relatively constant in time (on a short-term scale).
Most of the isotopes with atomic mass numbers below 14 decay to isotopes of carbon, while most of the isotopes with masses above 15 decay to isotopes of oxygen. The shortest-lived known isotope is nitrogen-10, with a half-life of 143(36) yoctoseconds , though the half-life of nitrogen-9 has not been measured exactly.