Atomic weight

The atomic mass of a chemical element is the mass of an atom at rest, most often expressed in unified atomic mass units. The atomic mass is often synonymous with relative atomic mass, average atomic mass and atomic weight; however, it is subtly different in that it can either be the abundance-weighted average of isotope masses of an element or the mass of a single isotope.

The relative atomic mass (also known as atomic weight and average atomic mass) is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance. Periodic tables usually list these with reference to the local environment of Earth's crust and atmosphere. An uncertainty in parenthesis is often included. For artificial elements the nucleon count of the most stable isotope is listed in brackets as the atomic mass.

The relative isotopic mass is the relative mass of the isotope, scaled with carbon-12 as exactly 12. No other isotopes have whole number masses due to the different mass of neutrons and protons, as well as loss/gain of mass to binding energy. However, since mass defect due to binding energy is minimal compared to the mass of a nucleon, rounding the atomic mass of an isotope tells you the total nucleon count. Neutron count can then be derived by subtracting the atomic number.

The pattern in the amounts the atomic masses deviate from their mass numbers 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 following: nuclear fission in an element heavier than iron produces energy, and fission in any element lighter than iron requires energy; the opposite is true of nuclear fusion reactions - fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

A similar definition applies to molecules; it is then called molecular mass. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms multiplied by the ratios of elements given in the chemical formula. A similar formula mass can be calculated for those compounds which do not form molecules.

Direct comparison and measurement of the masses of atoms and molecules is achieved with mass spectrometry.

The standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are 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 6.02 × 1023. One mole of a substance always contains almost exactly the atomic or molecular mass of that substance (which is the concept of molar mass), expressed in grams. For example, the atomic mass of iron is 55.847, and therefore one mole of iron has a mass of 55.847 grams. The formulaic conversion between atomic mass and SI mass in grams for a single atom is

where amu is the atomic mass unit and NA is Avogadro's number.

N.B: The generally accepted A-Level Chemistry definition of relative atomic mass is: "The average mass of an atom in a normal sample of an element expressed in terms of the mass of an atom of the most abundant isotope of carbon, taken as 12.0000"